homogeneous nickel and iron complexes for carbon dioxide
TRANSCRIPT
Homogeneous Nickel and Iron Complexes for Carbon Dioxide and Methanol
Transformations: Progress toward Commercial Chemical Production
By
Elizabeth M. Lane
B.A., Dartmouth College 2008
M.A., Brown University 2015
A Dissertation Submitted in Partial Fulfillment of the Requirements for the Degree of
Doctor of Philosophy in the Department of Chemistry at Brown University
Providence, Rhode Island
May 2018
© Copyright 2018 by Elizabeth M. Lane
iii
This dissertation by Elizabeth M. Lane is accepted in its present form by the Department
of Chemistry as satisfying the dissertation requirement for the degree of Doctor of
Philosophy.
Date: ________________ ____________________________
Wesley H. Bernskoetter, Advisor
Recommend to the Graduate Council:
Date: _________________ ____________________________
Eunsuk Kim, Reader
Date: _________________ ____________________________
Paul Williard, Reader
Approved by the Graduate Council:
Date: _________________ _________________________________________
Andrew G. Campbell, Dean of the Graduate School
iv
Curriculum Vitae
Elizabeth M. Lane graduated from Mohawk Trail Regional High School in 2004 as
valedictorian of her class and attended Dartmouth College in Hanover, New Hampshire for
her undergraduate studies. At the end of her junior year, she received an Arnold and Mabel
Beckman Research Scholarship and began undergraduate chemistry research in the
laboratory of Professor David Glueck, where she pursued the synthesis and structural
characterization of novel gold-phosphido complexes and completed an honors thesis. She
graduated magna cum laude from Dartmouth in 2008 with a Bachelor of Arts in Chemistry
and Philosophy. She then attended the University of Chicago and worked under the tutelage
of Professor Gregory Hillhouse in the field of organometallic chemistry, and received a
Master of Science in 2009. She spent several years as a high school chemistry teacher and
as an undergraduate laboratory manager, and then entered the doctoral program at Brown
University in the fall of 2013. She joined the laboratory of Professor Wesley Bernskoetter,
where she has studied the activation and transformation of carbon dioxide and methanol
using complexes containing earth-abundant metals like nickel and iron.
v
Acknowledgements
I feel very fortunate to have completed my graduate studies in the chemistry
department at Brown. The dedication and enthusiasm with which the faculty and staff
approach their duties is apparent, and I will not forget the sense of inclusion that I
experienced when working here. It is definitely an environment that fosters collaboration
between disciplines that is difficult to find elsewhere.
I would first very much like to thank my boss, Wes, for accepting me into his lab
and letting me to pretend like I owned it for five years. His deep understanding of chemistry
is something that I have always admired, and the clarity and patience with which he
conveys that knowledge to others is a rare skill. I cannot fully express how much I value
the time I have spent working for him. His lab has been a home to me, and for that I will
always be grateful.
I have also been gifted a second amazing teacher in our collaborator, Professor
Nilay Hazari, who I would like to thank for his insightful discussions as well as his frequent
encouragement. In addition, I would like to extend my sincere thanks to my committee
members, Professor Eunsuk Kim and Professor Paul Williard, for their years of advice and
their continued support.
I would also like to thank my labmates, a term which I deem to be inclusive for life.
Both the lab group I entered into as well as the one I’m leaving were full of intelligent,
hardworking, fun people who made the hard days bearable and the good days remarkable.
Thank you all, so much, for getting me through this. A special thanks to Alex, Ryan, and
Kevin, who have been like brothers to me, and to Yuanyuan, Ariana, Zoe, Melissa, and
Katie for their friendship and for many days spent commiserating.
vi
A final but equally important group of people to acknowledge is my family. Thank
you Mom, Dad, and Kate, for everything that you did to help get me here. Your strength,
intelligence, and kindness are what I aspired toward when I was growing up. You supported
me, loved me, and moved me back and forth across the country more times than can be
rationally expected of any one group of people. I absolutely could not have done this
without you.
vii
Table of Contents
Vita…………………………………………………………………………………iv
Acknowledgements………………………………………………………………...v-vi
List of Figures……………………………………………………………………...x-xi
List of Schemes…………………………………………………………………….xii-xiv
List of Tables……………………………………………………………………….xv
List of Abbreviations……………………………………………………………….xvi
General Introduction: The Significant Role of Homogeneous Catalysts in the Field of Small
Molecule Activation………………………………………………………………...1-6
Chapter 1: The Application of Homogeneous Nickel and Iron Complexes in the
Transformation of CO2 and CO2 Analogues into Valuable Organic Products……...7
1.1: Introduction: Carbon Dioxide as a Carbon Feedstock………………….8-16
1.2: Transition-Metal-Catalyzed Oxidative Coupling of CO2 and CO2 Analogues
with Olefins………………………………………………………………….17
1.2.1: Introduction………………………………………………..….17-27
1.2.A.i.: Nickel-Promoted CO2-Alkene Coupling for the Stereoselective
Production of Substituted Cyclic Anhydrides……………………….28
1.2.A.i.1: Introduction……………………………………….28-33
1.2.A.i.2: Results and Discussion…………………………...34-40
1.2.A.i.3: Summary……………………………………….…41
viii
1.2.A.ii.: The Use of CO Surrogates in Nickel-Promoted CO2-Alkene
Coupling for the Catalytic Production of Cyclic Anhydrides….…..42
1.2.A.ii.1: Introduction………………………………….….42-43
1.2.A.ii.2: Results and Discussion……………………….…44-47
1.2.A.ii.3: Summary…………………………………….…..48-49
1.2.B.: Nickel-Promoted Coupling of Isocyanates and Alkenes for the
Production of N-Substituted Acrylamides………………………….50
1.2.B.1: Introduction………………………………………..50-55
1.2B.2: Results and Discussion…………………………….55-65
1.2.B.3: Summary…………………………………………..65-66
1.2.C.: Section 1.2 Summary……………………………………….66-67
1.3: Transition-Metal-Catalyzed Hydrogenation of CO2 to Methanol…68
1.3.1: Introduction……………………………………….…..68-79
1.3.2: Results and Discussion……………………………….79-86
1.3.3: Summary……………………………………………...86-87
1.4: Chapter Summary………………………………………………88-89
1.5: References……………………………………………………....90-107
Chapter 2: The Application of an Iron Catalyst in the Transformation of Methanol into
Formamides and Ureas via Acceportless Dehydrogenation Reactions……………...108
ix
2.1: Introduction: Methanol as a Carbon Feedstock……………………….109-113
2.2: Iron-Catalyzed Amide Formation from the Dehydrogenative Coupling of
Alcohols and Secondary Amines…………………………………………..114
2.2.1: Introduction……………………………………………….…114-119
2.2.2.: Results and Discussion………………………………….......119-127
Catalytic Studies
Mechanistic Investigations
2.2.3: Summary…………………………………………………….127-128
2.3: Iron-Catalyzed Urea Synthesis: Dehydrogenative Coupling of Methanol and
Amines………………………………………………………………..……129
2.3.1: Introduction…………………………………………………129-133
2.3.2: Results and Discussion…………………………………...…133-142
Catalytic Studies
Mechanistic Investigations
2.3.3: Summary…………………………………………………….142-143
2.4: Chapter Summary……………………………………………………..144-145
2.5: References……………………………………………………………..146-155
Appendix 1…………………………………………………………………………156-228
Appendix 2…………………………………………………………………………229-272
x
List of Figures
Figure 1.1.1: CO2 reduction potentials at pH 7 (E° vs. NHE)……………………….…13
Figure 1.1.2: Metal binding modes of CO2………………………………………….....14
Figure 1.2.1: Depiction of β-H orientation in a nickelalactone complex and corresponding
d-orbital splitting diagram……………………………………………………………....24
Figure 1.2.A.1: Examples of [N,N], [P,N], and [P,P] types of bidentate ligands. tBu = tert-
butyl, Cy = cyclohexyl…………………………………………………………………..32
Figure 1.2.A.2: Structure key for ligand numbering. iPr = isopropyl, tBu = tert-butyl....34
Figure 1.2.A.3: Chiral nickel complexes competent for coupling propylene, CO2, and CO
to give MSA. iPr = isopropyl, tBu = tert-butyl………………………………………….40
Figure 1.2.A.4: CO surrogates and bases used, with abbreviations……………………44
Figure 1.2.B.1: 31P{1H} NMR spectrum in C6D6 of the nickel-promoted coupling products
from the ethylene/ethyl isocyanate reaction (2 atm/10 equiv.). Products C/D do not
incorporate ethylene. 300 MHz spectrometer…………………………………………...58
Figure 1.2.B.2: 31P{1H} NMR spectrum in C6D6 of the nickel-promoted coupling products
from the ethylene/tert-butyl isocyanate reaction (2 atm/10 equiv.) See Figure 1.2.B.1 for
labeling guide. 300 MHz spectrometer………………………………………………….60
Figure 1.2.B.3: 31P{1H} NMR spectrum in C6D6 of the nickel-promoted coupling products
from the ethylene/phenyl isocyanate reaction (1 atm/2 equiv.). See Figure 1.2.B.1 for
labeling guide. 300 MHz spectrometer…………………………………………………..62
Figure 1.2.B.4: Alternative ligands tested in Ligand-Ni(COD)-promoted
ethylene/isocyanate coupling reactions………………………………………………......64
xi
Figure 1.3.1: Previous (1-5) and current metal catalysts for CO2 hydrogenation to
methanol via CO2 derivatives. OTf = triflate, OAc = acetate…………………………..77
Figure 1.3.2: Structures of the iron-pincer complexes 1, 2, and 1-formate…………...79
Figure 2.2.1: A top-performing ruthenium catalyst (a) and all previous base-metal catalysts
(b) for dehydrogenative amide formation from amines and alcohols…………………. 116
Figure 2.2.2: Efficient ruthenium, manganese, and iron catalysts for the dehydrogenative
coupling of alcohols and 2° amines……………………………………………………..118
Figure 2.2.3: Variable temperature 31P NMR spectra indicate an equilibrium between 1
and 2 in the presence of 6 equiv. of methanol. Sample dissolved in benzene-d6……….126
Figure 2.3.1: Ruthenium and iron catalysts for symmetric and unsymmetric urea formation
from the dehydrogenative coupling of methanol and amines…………………………..132
xii
List of Schemes
Scheme 1.1.1: Industrial-scale syntheses incorporating CO2 as a carbon source…….….10
Scheme 1.1.2: Select C-O, C-N, C-C, and C-H bond-forming reactions with CO2. Products
in blue are industrially produced……………………………………………………...…..15
Scheme 1.2.1: Generic metal-catalyzed oxidative coupling between CO2 and an
unsaturated hydrocarbon (without a formally metal-bound CO2)……………………….18
Scheme 1.2.2: Commercial acrylic acid production…………………………………….21
Scheme 1.2.3: Proposed catalytic cycle for the transition-metal-catalyzed production of
acrylic acid from ethylene and CO2……………………………………………………..22
Scheme 1.2.4: Examples of nickel- and molybdenum-promoted oxidative coupling
reactions between ethylene and CO2…………………………………………………....23
Scheme 1.2.5: Methods of inducing acrylate removal from nickelalactones using (A) base
or (B) Lewis acids (LA)………………………………………………………………...26
Scheme 1.2.A.1: Proposed mechanism for the nickel-catalyzed production of cyclic
anhydrides from CO2, alkenes, and CO surrogates…………………………………….28
Scheme 1.2.A.2: Most prevalent methods of cyclic anhydride production…………... 30
Scheme 1.2.A.3: (A) Achiral and (B) chiral nickel-promoted methods for the production
of cyclic anhydrides…………………………………………………………………....31
Scheme 1.2.A.4: Methods of nickelalactone production from (A) CO2/alkene coupling and
retrosynthetically from (B) a cyclic anhydride or (C) acrylic acid…………………….35
Scheme 1.2.A.5: Lewis acid enhancement of the 3-,4-, and 5-Ni(COD) coupling reactions
between ethylene and CO2……………………………………………………………..38
xiii
Scheme 1.2.A.6: Previous application of CO surrogates with nickel species for the
production of cyclic lactams and the current proposed use for releasing cyclic
anhydrides…………………………………………………………………………..…...43
Scheme 1.2.B.1: Large-scale methods of synthesizing acrylamides…………………...51
Scheme 1.2.B.2: Nickel-catalyzed isocyanate/alkene coupling reactions with conditions
for accessing specific acrylamide regioisomers as well as propionamides and
succinimides…………………………………………………………………………….54
Scheme 1.2.B.3: Nickel-catalyzed isocyanate/alkene coupling reactions pursued in this
work…………………………………………………………………………………….55
Scheme 1.3.1: Products of CO2 hydrogenation reaction and the Gibbs free energy for
formic acid and methanol generation…………………………………………………..68
Scheme 1.3.2: Industrial method of methanol production from syngas, depicting (1) the
component equilibria and (2) the overall reaction……………………………………..69
Scheme 1.3.3: General method of H2 storage using reversible metal-catalyzed CO2
hydrogenation/dehydrogenation………………………………………………………..71
Scheme 1.3.4: Methods of indirect hydrogenation of CO2 to methanol……………….73
Scheme 1.3.5: Metal-catalyzed CO2 hydrogenation to methanol via (1) esters or (2)
formamides……………………………………………………………………………...74
Scheme 1.3.6: Flow diagram for the one-pot, two-step reaction setup………………....81
Scheme 2.1.1: (1) Industrial methanol production from syngas and (2) industrial
formaldehyde production from methanol……………………………………………….109
Scheme 2.1.2: Some of the major chemical species derived from methanol and their
uses……………………………………………………………………………………...110
xiv
Scheme 2.2.1: Metal-free methods for amide production……………………………...114
Scheme 2.2.2: Metal-promoted pathways for dehydrogenative amidation…………….115
Scheme 2.2.3: Iron-catalyzed coupling of benzaldehyde or methyl formate with
morpholine……………………………………………………………………………..125
Scheme 2.3.1: Metal-free pathways for urea formation……………………………….129
Scheme 2.3.2: Metal-catalyzed urea synthesis from non-methanol sources…………..130
Scheme 2.3.3: Metal-promoted pathway for dehydrogenative urea formation from
methanol and amines……………………………………………………………….…..131
Scheme 2.3.4: Potential scrambling pathways in unsymmetric urea formation……….138
Scheme 2.3.5: Proposed mechanistic pathways for metal-catalyzed urea formation starting
from methanol and primary amines……………………………… …………………..141
Scheme 2.3.6: Formamide-dependent synthesis of 1,1,3-tripentylurea……………..…141
xv
List of Tables
Table 1.2.A.1: Reaction of dcpe-Ni-lactone and dcpe-Ni-lactone’ Species with CO
Surrogates……………………………………………………………………………..46
Table 1.3.1: Select Reaction Optimization Data for CO2 Hydrogenation to Methanol in the
Presence of Morpholine……………………………………………………………….84
Table 2.2.1: Dehydrogenative Amidation of Methanol Catalyzed by 1……………...121
Table 2.2.2: Alcohol Dehydrogenation in the Presence of Morpholine Catalyzed by
1………………………………………………………………………………………..123
Table 2.3.1: Dehydrogenative Symmetric Urea Formation from Methanol and Amines
Catalyzed by 1………………………………………………………………………....134
Table 2.3.2: Dehydrogenative Unsymmetric Urea Formation from Formamides and
Amines Catalyzed by 1………………………………………………………………...137
xvi
List of Abbreviations
iPr = isopropyl LA = Lewis acid
tBu = tert-butyl SA = succinic anhydride
Cy = cyclohexyl MSA = methylsuccinic anhydride
Ph = phenyl NaBAr4F = sodium tetrakis
Et = ethyl [3,5-bis(trifluoromethyl)phenyl]borate
Me = methyl DBU = 1,8-diazabicyclo[5.4.0]undec-
CO2 = carbon dioxide 7-ene
CO = carbon monoxide NaOtBu = sodium tert-butoxide
H2 = dihydrogen LiOTf = lithium
COD = 1,5-cyclooctadiene trifluoromethanesulfonate
cdt = 1,5,9-cyclododecatriene Complex 1 = (iPrPNP)Fe(H)(CO)
dcpe = 1,2-bis(dicyclohexylphosphino)ethane iPrPNP = N[CH2CH2(PiPr2)]2
TON = turnover number
TOF = turnover frequency
Mt = million metric tons
kt = kilotons
ppm = parts per million
syngas = synthesis gas
[CO] = CO surrogate
[B] = base
NMR = nuclear magnetic resonance
GCMS = gas chromatography-mass spectrometry, FID = flame ionization detector
1
General Introduction:
The Significant Role of Homogeneous Catalysts in the Field of Small Molecule
Activation
2
The overwhelming dependence on fossil fuels for supplying the world’s energy
needs is unsustainable, and the rapid depletion of these non-renewable resources has
motivated investigations into alternative means of meeting global demands.1-3 Small
molecules such as dioxygen (O2), dinitrogen (N2), dihydrogen (H2), carbon dioxide (CO2),
and methane (CH4) appear in various biological metabolic processes and large-scale
industrial conversions, exhibiting their participation in the efficient and reversible storage
of energy in the form of chemical bonds.4 This combined with their abundance makes them
attractive targets for the development of alternative chemical feedstocks and fuels.4-7,9,13
Reactions of interest include the conversion of N2 to ammonia or hydrazine,8,9 the
reversible (de)hydrogenation of CO2 as a means of H2 storage for alternative fuels,10 H2
oxidation/O2 reduction coupled fuel cells,11 and CH4 (alkane) functionalization to higher
chemicals.12 However, a major obstacle to realizing these objectives is the thermodynamic
and often also kinetic stability of these materials.1,13 Those compounds that are more
amenable kinetically often suffer from a lack of product control instead (e.g. over-oxidizing
a substrate using O2).13 Addressing these problems is where homogeneous catalysts are
poised to contribute significantly to the advancement of the field of small molecule
activation. In addition to alleviating kinetic barriers, catalysts can enable otherwise
unavailable interactions between the more inert species listed above and highly reactive
molecules to yield a product with a more favorable thermodynamic profile.14-16 They can
also provide a means of controlling reactivity.13 It should be noted that currently the most
active catalysts for enacting these types of transformations are heterogeneous.8,17 Areas
where homogeneous catalysts gain an advantage are in enhanced product selectivity and a
greater mechanistic understanding that allows for the rational tuning of catalytic
3
behavior.8,18-20 The latter characteristic is especially crucial when developing catalytic
systems for new processes. A host of homogeneous catalysts have therefore been employed
in the task of facilitating small molecule transformations into value added products and
fuels. The application of catalysts containing inexpensive and earth-abundant metals
toward achieving this goal is of particular importance from both a sustainability and an
economic viewpoint.21 The work presented here details several such examples, using
homogeneous nickel and iron complexes for the functionalization of CO2 and methanol,
with the ultimate goal of catalyzing the production of commercially valuable chemicals
from cheap, widely available, and potentially renewable C1 carbon sources.
4
A note concerning catalyst comparisons: When considering catalytic performances, terms
such as conversion, yield, turnover frequency (TOF), and turnover number (TON) are often
used for comparisons between catalytic species. Conversion usually refers to the
consumption of one of the starting materials as a way of assessing reaction progress, but
this can occasionally be misleading if the selectivity of the reaction is poor. Yield is
determined based on the amount of the desired product in comparison to the theoretical
maximum and therefore avoids selectivity as a factor. While yield is a useful measure of
synthetic practicality, it often fails to express the full activity of a catalyst due to the ceiling
created by the limiting reagent(s). TOF explicitly references the performance of the catalyst
in question, analyzing how many molecules of product are generated per molecule of
catalyst over a given time period. The TOF, like most reaction rates, will change over the
course of catalysis, and is typically reported as the ability of a catalyst to deliver some
number of moles of product in one hour. This can be a valuable method to compare the
nativity of catalysis, but care must be taken to compare catalysts using TOFs determined
over the same reaction time periods. TON is a measurement of the total amount of product
delivered over the whole of the reaction per unit of catalyst. Several factors impact the
maximum TON of a catalyst, including the TOF, reaction selectivity, and catalyst lifetime.
When catalysis is conducted in a large excess of substrate, the maximum TON represents
the greatest number of times a molecule of catalyst can perform the desired transformation
before it becomes inactive. When developing new catalysts, we have selected TON as the
primary point of comparison because it is a value that cannot be easily skewed by changing
experimental methods and reflects several aspects important to catalysis in a single value.
5
References
1 Olah, G.A.; Goeppert, A.; Prakash, G.K.S. “Chemical Recycling of Carbon Dioxide to Methanol and
Dimethyl Ether: From Greenhouse Gas to Renewable, Environmentally Carbon Neutral Fuels and Synthetic
Hydrocarbons,” J. Org. Chem., 2009, 74, 487-498.
2 Appel, A.M.; Bercaw, J.E.; Bocarsly, A.B.; Dobbek, H.; DuBois, D.L.; Dupuis, M.; Ferry, J.G.; Fujita, E.;
Hille, R.; Kenis, P.J.A.; Kerfeld, C.A.; Morris, R.H.; Peden, C.H.F.; Portis, A.R.; Ragsdale, S.W.; Rauchfuss,
T.B.; Reek, J.N.H.; Seefeldt, L.C.; Thauer, R.K.; Waldrop, G.L. “Frontiers, Opportunities, and Challenges
in Biochemical and Chemical Catalysis of CO2 Fixation,” Chem. Rev., 2013, 113, 6621-6658.
3 Goeppert, A.; Czaun, M.; Jones, J.-P.; Prakash, G.K.S.; Olah, G. “Recycling of carbon dioxide to methanol
and derived products- closing the loop,” Chem. Soc. Rev., 2014, 43, 7995-8048.
4 Meyer, F.; Tolman, W.B. “Forums on Small-Molecule Activation: from Biological Principles to Energy
Applications,” Inorg. Chem., 2015, 54, 5039-5039.
5 Lehnert, N.; Peters, .J.C. “Preface for Small-Molecule Activation: from Biological Principles to Energy
Applications. Part 2: Small Molecules Related to the Global Nitrogen Cycle,” Inorg. Chem., 2015, 54, 9229-
9233.
6 Darensbourg, M.Y.; Llobet, A. “Preface for Small-Molecule Activation: from Biological Principles to
Energy Applications. Part 3: Small Molecules Related to (Artificial) Photosynthesis,” Inorg. Chem., 2016,
55, 371-377.
7 Fujita, E.; Goldman, A.S. “Preface for Small-Molecule Activation: Carbon-Containing Fuels,” Inorg.
Chem., 2015, 54, 5040-5042.
8 Nishibayashi, Y. “Recent Progress in Transition-Metal-Catalyzed Reduction of Molecular Dinitrogen under
Ambient Reaction Conditions,” Inorg. Chem., 2015, 54, 9234-9247.
9 Bhattacharya, P.; Prokopchuk, D.E.; Mock, M.T. “Exploring the role of pendant amines in transition metal
complexes for the reduction of N2 to hydrazine and ammonia,” Coord. Chem. Rev., 2017, 334, 67-83.
10 Wang, W.-H.; Himeda, Y.; Muckerman, J.T.; Manbeck, G.F.; Fujita, E. “CO2 Hydrogenation to Formate
and Methanol as an Alternative to Photo- and Electrochemical CO2 Reduction,” Chem. Rev., 2015, 115,
12936-12973.
11 Zhang, W.; Lai, W.; Cao, R. “Energy-Related Small Molecule Activation Reactions: Oxygen Reduction
and Hydrogen and Oxygen Evolution Reactions Catalyzed by Porphyrin- and Corrole-Based Systems,”
Chem. Rev., 2017, 117, 3717-3797.
12 Munz, D.; Strassner, T. “Alkane C-H Functionalization and Oxidation with Molecular Oxygen,” Inorg.
Chem., 2015, 54, 5043-5052.
13 Zimmermann, P.; Limberg, C. “Activation of Small Molecules at Nickel(I) Moieties,” J. Am. Chem. Soc.,
2017, 139, 4233-4242.
14 Huang, K.; Sun, C.-L.; Shi, Z.-J. “Transition-metal-catalyzed C-C bond formation through the fixation of
carbon dioxide,” Chem. Soc. Rev., 2011, 40, 2435-2452.
6
15 Pearson, R.L.; Eisaman, M.D.; Turner, J.W.G.; Edwards, P.P.; Jiang, Z.; Kuznetsov, V.L.; Littau, K.A.;
Di Marco, L.; Taylor, S.R.G. “Energy Storage via Carbon-Neutral Fuels Made from CO2, Water, and
Renewable Energy,” Proceedings of the IEEE, 2012, 100, 440-460.
16 Boddien, A.; Gärtner, F.; Federsel, C.; Piras, I.; Junge, H.; Jackstell, R.; Beller, M., Catalytic Utilization
of Carbon Dioxide: Actual Status and Perspectives, in Organic Chemistry: Breakthroughs and Perspectives;
Ding, K., Dai, L.-X., Eds.; Wiley-VCH: Weinheim, Germany, 2012, pp. 685-724.
17 Burford, R.J.; Yeo, A.; Fryzuk, M.D. “Dinitrogen activation by group 4 and group 5 metal complexes
supported by phosphine-amido containing ligand manifolds,” Coord. Chem. Rev., 2017, 334, 84-99.
18 Huff, C.A.; Sanford, M.S. “Cascade Catalysis for the Homogeneous Hydrogenation of CO2 to Methanol,”
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7
Chapter 1:
The Application of Homogeneous Nickel and Iron Complexes in the Transformation of
CO2 and CO2 Analogues into Valuable Organic Products
8
1.1 Introduction: Carbon Dioxide as a Carbon Feedstock
The most recent Renewables Global Status Report indicates that in 2015, 78.4% of
global energy consumption stemmed from fossil fuels, with renewable sources (such as
wind, solar, geothermal, and biomass) accounting for only 19.3%, and the remaining 2.3%
belonging to nuclear power.1 This worldwide dependence on non-renewable energy
resources has not only had significant deleterious environmental effects,2 but faces an
inevitable future where these finite resources are gone and will not be realized again on a
human timescale.3,4 One consequence of fossil-fuel dependence is that the cost of these
resources will increase exponentially as their availability plummets.4 Another is the ever-
increasing production of carbon dioxide (CO2) due to the burning of fossil fuels, the rising
atmospheric levels of which have been implicated in widespread destructive environmental
alterations, including climate change and ocean acidification.4,5 Atmospheric CO2 content
is dramatically higher than its pre-industrial level (from 250 ppm to 410 ppm in 2017),6,7
is increasing at a rate of ~1% per year,5 and despite efforts to decrease the contribution
from energy consumption,8 the natural forms of CO2 sequestration such as plant
photosynthesis and aqueous sinks5 cannot keep up. There are therefore significant
economic and environmental motivations to transition to renewable methods of meeting
the global fuel demand for energy, transportation, and commodity chemical applications.
One such method for partially relieving CO2 concentrations while decreasing fossil fuel
dependence is to utilize CO2 as a carbon feedstock in the production of alternative fuels
and commercial chemicals.6
CO2 is an attractive C1 carbon source for organic synthesis because it is cheap,
abundant, renewable, and non-toxic (in comparison to other sources like phosgene or
9
carbon monoxide (CO)).6,9-13 As opposed to techniques aimed only at carbon capture and
storage, carbon recycling methods both remove CO2 from the atmosphere and transform it
into practically useful species while avoiding issues of high transportation costs to storage
sites and dependability of storage time periods.14,15 It is also the only method that generates
a profit from what is otherwise a waste item.14 While the utilization of CO2 for fuels and
commodity chemical production is unlikely to significantly counteract anthropogenic
emissions,10 it is an economically beneficial way of sequestering CO2 for multiple
lifetimes. By replacing organic precursors typically derived from non-renewable sources
with inexpensive CO2, the energy and cost of chemical production is reduced (with a
corresponding price reduction for the consumer) while also decreasing fossil fuel
dependency.10,15-18 If this synthesis of functionalized desirable species from CO2 is
accomplished using energy sources like solar, wind, or geothermal, it serves as a possible
avenue for the storage and transportation of renewable energy in chemical form, without
necessitating a change in the energy infrastructure.15,18 Alternatively, CO2-derived fuels
themselves can be viewed as renewable, carbon-neutral energy sources,3,15,18 so their
synthesis and subsequent use pulls current CO2 from circulation while also diverting from
future emissions by substituting for fossil fuels.3
Evidence for the feasibility of CO2 conversion is found in biological systems, which
over billions of years have evolved enzymatic pathways for processing CO2 (via the
making of C-H/C-C bonds and the breaking of C-O bonds) and thereby effectively storing
and utilizing energy in the form of chemical bonds.4 Despite the aforementioned
motivations for using CO2 in the production of high-value chemicals as well as the plethora
of examples of its application in living systems, the actual large-scale use of CO2 as a
10
carbon source is quite limited. Currently, only ~1% of CO2 produced is used for chemical
synthesis,5,10 and the few CO2-fixation products manufactured on an industrial scale are
carbonates (inorganic, cyclic, and polymeric), urea, salicylic acid, and (to a lesser extent)
methanol (Scheme 1.1.1).6,10
Carbonates:
Inorganic carbonates have applications in the production of paper, plastics, rubber,
paint, and building materials, and calcium carbonate (CaCO3) alone accounts for 15 Mt
Scheme 1.1.1: Industrial-scale syntheses incorporating CO2 as a carbon source.
11
(million metric tons) of the overall global market of more than 200 Mt per year (not entirely
derived from industrial CO2 sources).15,19 Therefore, the use of industrial residues rich in
magnesium, calcium, or sodium in conjunction with emitted CO2 to give carbonates has
the potential to be very lucrative. In fact, Skymine has received over $25 million in funding
to apply a patented CO2 mineralization process by which it converts flue gas CO2 into
sodium bicarbonate.15 Cyclic carbonates are made commercially from epoxides and CO2
(Scheme 1.1.1, reaction 1),20 but methods of replacing the pre-synthesized epoxides with
the direct reaction of alkenes and CO2 have been developed.21 Their value is largely based
on their role in the production of linear carbonates and as monomers for polycarbonates,
the combined market for which amounts to approximately 20 Mt per year. Uses include
resins for data storage, organic reagents for alkylation and acylation reactions, solvents
(e.g. for pharmaceuticals preparations), and battery electrolytes.10,14,15,20,22
Urea
The commercial production of urea from ammonia and CO2 via a carbamate
intermediate (Scheme 1.1.1, reaction 2) is known as the Bosch-Meiser urea process and
has been in use since the 1920s.15,19 It is conducted under elevated temperature (~160 °C)
and pressure (~110 atm) conditions, where the exothermic first step drives the endothermic
second step. The >150 Mt per year demand for urea represents the largest-volume industrial
example of converting CO2 to a value-added product, with its primary uses being as a
fertilizer and resin precursor.15 The use of homogeneous catalysts with substituted amines
and CO2 has yielded functionalized ureas but has not yet been realized on a commercial
scale (see Chapter 2, section 2 for more information).
Salicylic Acid
12
The synthesis of salicylic acid was developed in the late 1800’s and is shown in
Scheme 1.1.1, reaction 3 (o-isomer is shown but the reaction also produces the p-isomer).
It is also known as the Kolbe-Schmitt reaction and is performed at 120-130 °C and ~100
atm CO2.23 Salicylic acid is used as a food preservative and in a variety of pharmaceuticals,
including those for the treatment of dermatological conditions (acne, psoriasis, ringworm,
dandruff) and for pain relief. The latter accounts for its most widespread application, which
is in the production of aspirin on a 20 kt (kilotons) per year scale.19,24
Methanol
Methanol is produced on a more than 80 Mt industrial scale from a synthesis gas
mixture (syngas; H2/CO2/CO) using heterogeneous alumina-supported mixed copper and
zinc oxide catalysts at elevated temperatures and pressures (50-100 atm, 200-300 °C;
Scheme 1.1.1, reaction 4).25-28 While a small percentage of manufactured methanol is used
as a solvent, most serves as a reagent in the synthesis of a variety of valuable end-products
like textiles, plastics, film, paints, adhesives, and coatings.29 This topic is covered in greater
detail in the last section of this chapter as well as in Chapter 2.
These limited examples offer the possibility of a renewable CO2-based carbon
economy, however, significant obstacles to the more widespread utilization of CO2 as a
carbon building block in the synthesis of organic molecules are its kinetic and
thermodynamic stabilities.5,6,18,30 Containing the most oxidized state of carbon,9 CO2 is
highly unreactive,6 making it a poor ligand and difficult to transform via direct reduction.31
Despite these complications, significant effort has been invested in achieving the
electrochemical reduction of CO2 to CO, methane, formic acid, or methanol.44 The appeal
stems not only from the recycling of CO2 and the production of valuable compounds, but
13
also from the storage of excess electrical energy in chemical form, as their energy density
can be anywhere from 10 to 100 times that of a battery.19,32 A one-electron reduction of
CO2 is highly unfavorable, and as a result proton-coupled multi-electron steps to yield the
products mentioned above are preferred (Figure 1.1.1).32,44 The electrocatalytic reduction
of CO2 has been studied since the 1980’s with either solid metal electrodes33 or
homogeneous transition metal complexes,32,34-37,44 however, there are significant
challenges with respect to low activities and the generation of side products like CO and
H2.32 Ideal catalysts employed in this process would increase overall activity and
selectivity, but would balance these kinetic advantages with a good match for the redox
potential of the electron-transfer reaction.32,44 These catalysts must also be capable of two-
electron transfer processes and be water stable (due to its use as a proton source). During
the electrochemical reduction reaction, several reductive equivalents build up at the
catalyst site, which then uses multi-electron pathways (hydride transfer, atom transfer,
insertion) to transform CO2.32 The most successful catalysts in this respect are iridium-
based,32,38,39 however, they (like other catalysts for this process) suffer from a trade-off
between the coveted traits of high activity, high selectivity, and low overpotential,32 and
usually sacrifice one for the other two. Photochemical reduction systems face similar
Figure 1.1.1: CO2 reduction potentials at pH 7 (E° vs. NHE).
14
issues, along with the added challenge of devising ways of enacting a multi-electron
reaction from a single photon, single-electron-transfer process.236 This in combination with
the incompatibility of the reductive conditions with many desirable products has motivated
the investigation of alternative means for CO2 transformations.
Although the energetic barriers for CO2 activation are significant, metal-mediated
non-electrochemical strategies for its incorporation into organic syntheses are promising
targets. A weakly Lewis acidic carbon and weakly Lewis basic oxygens due to bond
polarity make it possible for CO2 to bind to a metal center,4,10 and subsequent reduction
and transformation is much more favorable.42 However, not all metals have the basicity
necessary to accomplish this and metal-CO2 complexes are often unstable. Therefore, while
multiple examples of these complexes exist (the various binding modes of which are
depicted in Figure 1.1.2),10,17,40-43 it would be advantageous to use metal systems where
CO2 pre-coordination is not required for reaction but where participation in coupling and
insertion reactions is still possible. Current strategies aimed at CO2 activation combine CO2
with highly reactive starting materials (such as epoxides, dihydrogen, or unsaturated
compounds) to yield target species that are much lower in energy, thereby providing a
thermodynamic driving force and achieving a much more favorable thermodynamic
profile.9,18,44 The addition of a transition metal catalyst to mitigate unfavorable kinetic
Figure 1.1.2: Metal binding modes of CO2.
15
parameters5 that can also act on CO2 without it being formally metal-bound45 would
address any further resistance to functionalization.
Despite its chemical inertness, the slightly electrophilic C allows for the reaction of
CO2 with strongly nucleophilic reagents to give new C-H, C-O, C-N, and C-C bonds
(Scheme 1.1.2).30,44 Practical examples that rely on this principle in the large-scale
production of organic compounds are the previously mentioned cases of cyclic and
polymeric carbonates, urea, and salicylic acid.30,44 While reaction with energy-rich
substrates addresses some of the thermodynamic concerns surrounding CO2 utilization,
kinetic challenges can still remain.44 In order to alleviate these problems, the harsh reaction
conditions (high temperatures and pressures) demonstrated in the industrial examples
above and/or the use of metal catalysts are required. For example, even though the metal-
free production of cyclic carbonates has been commercialized for over 50 years, recent
investigations into the use of metal porphyrin catalysts containing aluminum, magnesium,
Scheme 1.1.2: Select C-O, C-N, C-C, and C-H bond-forming reactions with CO2.
Products in blue are industrially produced.
16
cobalt, chromium, iron, and zinc have significantly increased the reaction turnover under
milder conditions.30 The metal-free and metal-catalyzed manufacturing of new C-O and C-
N bonds using CO2 as a carbon feedstock has been extensively covered by several excellent
reviews and a book, and are beyond the scope of this discussion. They similarly cover
metal-free C-C bond formation using organolithium, organozinc, and Grignard reagents as
highly reactive and strong nucleophiles for the activation of CO2, as well as the metal-free
and metal-catalyzed reaction of less nucleophilic organoboron and organohalide
compounds with CO2, all giving carboxylic acids and their derivatives.9,30,44,46,47 While
these are diverse and effective ways of generating new C-C bonds incorporating CO2, they
are often highly intolerant of other functional groups. This chapter will instead focus on
the formation of new C-C and C-H bonds (inherently more difficult that C-N or C-O due
to decreased reagent nucleophilicity)30 using two metal-catalyzed CO2 activation methods:
the oxidative coupling of CO2 with olefins and the hydrogenation of CO2. These methods
for the transformation of CO2 (and CO2 analogues) will be discussed in the context of the
valuable chemicals they can be used to produce: substituted cyclic anhydrides and N-
substituted acrylamides with the former, and methanol with the latter.
17
1.2: Transition-Metal-Catalyzed Oxidative Coupling of CO2 and CO2 Analogues with
Olefins
1.2.1: Introduction
The first reported case of a metal-coordinated CO2 species by Aresta et al. in 1975
motivated extensive research focusing on the metal-catalyzed activation and
transformation of CO2.48 One such method involves oxidatively coupling CO2 (which is
highly stable) with reactive small molecules like alkenes or alkynes at a transition metal
center, whereby conversion to valuable organic molecules becomes more
thermodynamically favorable.49 The metal-promoted reaction of CO2 with unsaturated
hydrocarbons has been studied since the 1970’s, with pioneering work by Hoberg, Walther,
Inoue, and Yamamoto. Nearly all examples rely on the use of late, electron-rich transition
metals (d8-d10: Fe0, RhI, Ni0, Pd0 or PdII) with low oxidation states because they have
orbitals available for ligand backbonding, they can easily achieve higher oxidation states
(as part of the coupling process), and because of their nucleophilic nature. This allows for
a σ-type bonding interaction between the metal and the slightly electrophilic carbon of
CO2.10 Computational work by Pápai et al. has suggested that CO2 pre-coordination is not
necessary for these reactions and that coupling can occur in a concerted fashion,45 although
experimental evidence in some cases has indicated distinctly bound CO2 and alkene ligands
at the metal center.50 In cases where CO2 does not act as a discreet ligand, metal
nuclophilicity is still important because it facilitates alkene and alkyne binding via metal
to ligand backbonding, thereby activating the unsaturated species for subsequent CO2
coupling to form a metallocycle (Scheme 1.2.1).
18
Experimental investigation into the functionalization of CO2 using metal-promoted
coupling reactions can be assessed based on the CO2 co-reactant: alkenes, dienes, and
alkynes. Despite the diversity of substrates and over 30 years of research, there are very
few systems in which these transformations are catalytic.10 Some of the earliest work in
this area of CO2 activation was the palladium-catalyzed coupling reaction of CO2 and 1,3-
butadiene presented by Inoue et al. in 197651 and Musco et al. in 1978.52 These reactions,
however, yielded product mixtures where the desired carboxylic acid was a minor
contributor compared to multiple butadiene oligomers. These results are characteristic of
the behavior of dienes in general in this type of reaction since they tend to telomerize at
late transition metals, a trait which has been established for some time (e.g. the synthesis
of 1,5,9-cyclododecatriene, a.k.a. cdt, from butadiene and nickel metal).10,53-55 Some
progress in controlling product distribution has been made since the original publications
(such as by changing the metal starting material or using more sterically demanding
ligands),56,57 but it still remains the most significant challenge to using dienes as CO2
coupling partners. The metals capable of realizing this transformation are also quite
limited, and those that can do so catalytically primarily rely on palladium,51,52,56-60 with
only a few examples containing nickel46,61 and rhodium.62 The early work on the coupling
reaction between CO2 and alkynes faced similar selectivity issues with the formation of
dimers and trimers in addition to the desired pyrone product. Inoue et al. were also the first
Scheme 1.2.1: Generic metal-catalyzed oxidative coupling reaction between CO2 and
an unsaturated hydrocarbon (without a formally metal-bound CO2).
19
to publish on this topic, reacting a nickel (0) complex [Ni(COD)2, where COD = 1,5-
cyclooctadiene] with CO2 and 1-hexyne to yield a mixture of cyclotrimerization products,
only 10% of which was the desired species.63 Further development of this field using
rhodium64 and nickel catalysts65 eventually gave 100% yield with 96% selectivity for the
pyrone product using a sterically small phosphine ligand (triethylphosphine).66 However,
these results represent a very limited substrate scope and therefore have not been widely
applied. They also conflict with later findings that the selectivity of reactions between CO2
and diynes is improved with sterically larger ligands like N-heterocyclic carbenes.67,68
Overall, methods for controlling product preference in metal-promoted CO2/alkyne
coupling reactions are still poorly understood. In addition, many examples require the use
of aqueous acid or carbon monoxide (CO) to release the final product, obviating
catalysis.68-74 As a result, while coupling CO2 with dienes and alkynes is relevant to
accessing a range of functionalized and desirable products (covered in detail in the cited
reviews),6,9,10 the results presented in this chapter will instead focus on coupling reactions
of CO2 with alkenes.
The earliest work demonstrating CO2/olefin coupling was published by Inoue et al.
in 1979, featuring a palladium complex (Pd[dba]2, dba = dibenzylideneacetone) for the
production of lactones from CO2 and methylenecyclopropanes, which are highly reactive
due to pronounced ring strain. A mixture of two lactone products was observed, and the
identity of the majority species could be controlled by alteration of the ligand environment
on palladium.75 The original work did not address catalytic activity, and while future
investigations by other groups did, issues of side reactions between the product lactones
and starting material still in solution detracted significantly from yield.76 Therefore, while
20
these palladium results represent initial forays into the concept of metal-catalyzed coupling
of CO2 with a very restricted set of alkenes, it was investigations by the groups of Hoberg,
Walther, and Yamamoto in the 1980’s that established the scope of the field and identified
its most significant challenges.
Initial studies presented independently by Hoberg and Walther centered on the use
of nickel (0) sources containing weakly bound and easily displaced olefins like Ni(COD)2
and Ni(cdt) in the presence of bulky, electron-donating ligands like phosphines (e.g. 1,2-
bis(dicyclohexylphosphino)ethane = dcpe, triethylphosphine = PEt3) or bipyridine.77,82 The
larger ligand sizes discouraged binding multiple units to the metal center, which would
create catalytically inactive tetra-substituted species (NiL2 for bidentate ligands or NiL4 for
monodentate, both 18-electron species). The ligand σ-donating aspect contributed to the
overall nucleophilicity of the metal in order to promote metal-alkene binding and activation
toward further manipulation (via metal to ligand donation into the alkene antibonding
orbitals, weakening the carbon-carbon double bond) as well as to facilitate the necessary
increase in oxidation state upon CO2 coupling. These early works spanned alkenes as
simple as ethylene77-79 to sterically hindered species like norbornene,80 cyclopentene,81 and
dicyclopentadiene,82 as well as substituted alkenes78 and select linear dienes.83,84 Hoberg
et al. showed that under specific pressure conditions, nickel could facilitate the coupling
of multiple equivalents of an alkene (ethylene) with CO2.85 Hoberg’s group also extended
these concepts to include iron (0) complexes as metal promoters, some of which displayed
the unique characteristic of allowing multiple CO2 insertions to give diesters (after
methanolysis workup).79,84 It was demonstrated that while oxidative coupling of CO2 and
the selected alkene could easily proceed to the five-membered metallocycle in the presence
21
of a metal (known as nickelalactones for M = Ni), spontaneous elimination of the desired
product was not observed. While this permitted the extensive study of diverse metallocycle
structures (sometimes via ligand substitution),86 product liberation was usually attained by
acid hydrolysis or methanolysis (to yield a carboxylic acid or an ester, respectively) or by
addition of CO gas (to yield a cyclic anhydride).77,78,83 These methods, however, prevented
the catalytic application of CO2/olefin coupling toward synthesizing organic compounds
through decomposition of the metal species or conversion to an inactive metal carbonyl
complex (e.g. L2Ni(CO)2).
These pioneering works motivated extensive studies into the oxidative coupling of
CO2 with olefins and alkynes. Of particular interest was the coupling reaction between CO2
and ethylene because of the potential to liberate acrylic acid, and from these investigations
the major challenges associated with enacting these types of CO2 transformations have
been established. The growing market for acrylic acid was over 4.92 Mt in 2014.87 Its worth
is mostly dependent on its role as a precursor to polyacrylates, which are used to make
plastics, adhesives, coatings, detergents, and super absorbent polymers (such as those in
diapers).88 Current acrylic acid production depends on a two-step oxidation over
heterogeneous catalysts and begins with propylene (Scheme 1.2.2).89-91 Not only does this
process require harsh temperatures (250-400 °C)89,91 and go through a highly toxic acrolein
intermediate,89 but propylene is non-renewably sourced from fossil fuels and displays
Scheme 1.2.2: Commercial acrylic acid production.
22
higher costs as a result, which is highly correlated with the cost of acrylic acid.87 There is
therefore significant appeal in developing a method of acrylic acid production under milder
conditions and from cheaper, potentially renewable starting materials like ethylene and
CO2.10 A proposed catalytic cycle for this process is shown in Scheme 1.2.3. It begins with
binding the olefin to the metal, followed by oxidative coupling with CO2 to give the
metallocycle and then β-H elimination to the bound acrylate species. A final reductive
elimination step releases the product and re-opens a binding site for further catalysis.
Scheme 1.2.3: Proposed catalytic cycle for the transition-metal-catalyzed production
of acrylic acid from ethylene and CO2.
23
The groups of Hoberg and Carmona independently pursued the goal of metal-
catalyzed acrylic acid/acrylate formation from CO2/ethylene coupling using zerovalent
nickel and group VI metals, respectively (Scheme 1.2.4).77,78,92-95 Experimental and
computational mechanistic studies indicated that these metal complexes likely shared
common intermediates but faced different obstacles with respect to achieving
catalysis.45,50,96-100 In the case of complexes containing group VI metals (e.g. Mo, W), the
oxidative coupling of CO2 and ethylene occurred easily at ambient temperatures and
pressure, and β-H elimination to give the bound acrylate species was similarly facile.50,92-
95,101 However, the oxophilicity of group VI metals prevented product release via reductive
elimination in the absence of harsh bases (like butylithium)93 or strong electrophiles (such
as iodomethane),101 which are incompatible with catalysis.
Scheme 1.2.4: Examples of nickel- and molybdenum-promoted oxidative coupling
reactions between ethylene and CO2.
24
In contrast, acrylate elimination from nickel should occur more easily due to its
lower oxophilicity compared to metals like molybdenum and tungsten, but it faces greater
challenges in the coupling and β-H elimination stages of the catalytic cycle. Higher
temperatures and pressures are generally required to enact CO2/ethylene coupling using
nickel, which, while attainable, can cause prohibitive side reactions like ligand degradation
or multiple CO2 insertions.78,102,103,113 It is suggested that the ancillary ligand environment
can affect the kinetic barrier for coupling,45,104,113 but using it to control reaction energetics
is still not fully understood. By far the most significant issue with using nickel in
CO2/ethylene coupling, however, is its reticence toward β-H elimination, which is
attributed to two factors. The first is that crystallographic evidence indicates that the C-H
bonds on the lactone at the β position to the metal are oriented away from nickel and would
require significant ring distortion (and corresponding ring strain) to put them within
covalent bonding distance (Figure 1.2.1).10,78,86,105,106,109,111,113,149,151,152 The ease with
which group VI metals accomplish β-H elimination from similar metallocycles indicates
an additional factor, which is the necessity of an available empty metal orbital to accept
Figure 1.2.1: Depiction of β-H orientation in a nickelalactone complex and
corresponding d-orbital splitting diagram.
25
electron donation from the target C-H bond in order to form an agostic intermediate.50,92-
95,107,113 As shown in Figure 1.2.1, the empty dx2-y2 orbital in d8-square planar complexes
like nickelalactones is physically inaccessible for formation of the new Ni-H bond due to
the presence of the existing ligands.108,114
Solutions to this problem have focused on the use of additives to induce ligand
rearrangements that would re-orient the lactone β-hydrogens as well as open up access to
the necessary orbital. The first example of β-H elimination from a nickelalactone species
was achieved by Fischer et al. in 2006 through the addition of dppm (1,1-
bis(diphenylphosphino)methane) to give a dimeric Ni(II)-acrylate species.109
Unfortunately, this process was not catalytic and required the sacrifice of the dppm ligand,
so while it represented a significant step forward in understanding product release from
nickel, it would be a very expensive and wasteful way to generate acrylate. The groups of
Reiger and Kühn employed methyl iodide (MeI) to weaken the Ni-O bond, imparting more
flexibility into the lactone unit to allow for β-H elimination of methyl acrylate. Yields for
this reaction were low and the use of MeI as a reagent made catalysis impossible through
the formation of inactive Ni-halides species.110-112 Recent studies have focused on using an
external sodium base to deprotonate the lactone β-H directly (without the need for a transfer
to nickel) or incorporating a Lewis acid to facilitate ligand reorientation to allow for metal
participation, both with the subsequent release of sodium or lithium acrylate (Scheme
1.2.5).104 Limbach et al. initially applied strong bases like sodium tert-butoxide for this
purpose, and while good acrylate yields were realized, strong bases are incompatible with
the high CO2 pressures needed to initiate the coupling reaction. As a result, either isolation
of the initial nickelalactone prior to base addition or sequential additions of CO2, ethylene,
26
and base were needed to achieve even low TONs (~10), both of which are highly
impractical.113 Jin et al. developed this concept with the use of Lewis acids for lactone
activation and ring-opening, allowing for the incorporation of a milder base for
deprotonation to yield acrylate.114,115 This process, however, was still not catalytic.
Simultaneous publications in 2014 from the independent groups of Vogt and Limbach
detailed the first cases of one-step nickel-catalyzed CO2/ethylene coupling to produce
lithium or sodium acrylate, respectively. Vogt et al. focused on extending the applicability
of Lewis acids (e.g. LiI) through combination with a zinc reductant and mild bases (e.g.
Et3N) to release lithium acrylate with TONs <25.116 The base in this case was not for direct
deprotonation of the nickelalactone but for quenching the HI byproduct so that it would
not protonate their intended acrylate to propionate. Limbach’s work achieved higher
activities, with TONs of up to 107 for sodium acrylate production using sodium phenoxide
bases for deprotonation and zinc as a reducing agent, presumably to regenerate the
nickel(0) starting complex.117,118 While this accomplishment marks a milestone in the field
of using CO2 as a carbon feedstock for the generation of valuable functionalized
compounds, there are still significant limitations that must be overcome. For example, this
method requires multiple equivalents of a strong reductant, the turnover has been restricted
Scheme 1.2.5: Methods of inducing acrylate removal from nickelalactones using (A)
bases or (B) Lewis acids (LA).
27
by a side reaction of base with CO2 to form carbonates, and several grams of additives are
required for each reaction.
Rather than focusing on the development of these systems for inducing product
release, the β-H elimination issues can be circumvented altogether by shifting the focus
away from acrylates and toward other value-added chemicals like substituted cyclic
anhydrides or N-substituted acrylamides, which are the topics of the following sections. In
this way, the established reactivity of nickel with respect to CO2/olefin coupling reactions
can be exploited to generate a wider range of products. These behaviors have also been
extended to include systems involving species that are isoelectronic with CO2.
28
1.2.A.i.: Nickel-Promoted CO2-Alkene Coupling for the Stereoselective Production of
Substituted Cyclic Anhydrides
1.2.A.i.1: Introduction
Cyclic anhydrides (substituted or otherwise) can be produced from ring expansion
and functionalization of nickelalactone species via insertion of CO and subsequent
reductive elimination of the target (Scheme 1.2.A.1).77,153 In order to avoid the catalyst
deactivation pathway leading to inert L2Ni(CO)2 complexes, CO surrogates (denoted as
[CO]) could be employed as slow-release sources of CO. This subsection will address the
application of nickel complexes containing chiral ligands for the regio- and stereo-
Scheme 1.2.A.1: Proposed mechanism for the nickel-catalyzed production of cyclic
anhydrides from CO2, alkenes, and CO surrogates.
29
selective production of substituted cyclic anhydrides, while the following subsection
discusses using an achiral nickel complex for screening potential CO surrogates for
catalytic applications.
Cyclic anhydrides (such as succinic anhydride (SA) or methylsuccinic anhydride
(MSA)) are commercially valuable chemicals used in the production of
pharmaceuticals,119,120,134 agrochemicals,121 resins and coatings,121 flavors and
fragrances,122 plasticizers,123,124 and as reagents in organic synthesis.125-128 Those
containing a double bond in the five-membered ring, such as maleic anhydride, can
undergo Diels-Alder reactions with olefins to give alkenylsuccinic anhydrides, which are
in high demand in the paper industry129 and as petroleum additives.130,131 While the market
for cyclic anhydrides is highly product-dependent (e.g. 2.72 Mt in 1998 for phthalic
anhydride, 30,000 tons in 2008 for succinic anhydride)132,133 and individually much smaller
than that of acrylic acid and its derivatives,87 specialty applications of these organic
molecules makes them a desirable target. For example, making a substituted cyclic
anhydride like MSA enantioselectively is appealing from a pharmaceutical perspective,
where often only one enantiomer is biologically active or is effective for the intended
application.134
Cyclic anhydride production mainly relies on the metal-free (large-scale) or metal-
catalyzed (mostly research-based) dehydration of dicarboxylic acids (Scheme 1.2.A.2,
reaction 1)135-138 or the heterogeneously-catalyzed oxidation of benzene and benzene
derivatives (Scheme 1.2.A.2, reaction 2).139,140 Both operate under harsh temperature
conditions, require strong dehydrating or oxidizing agents, and are non-stereospecific.
Smaller-scale production from the metal-catalyzed carbonylation of alkynes,141,142 alkenoic
30
acids,143 and lactones144,145 has been investigated (Scheme 1.2.A.2, reactions 3-5), but these
methods proceed with low yields, significant side products, or limited substrate scope.
They are also not particularly stereoselective. Metal-catalyzed double carbonylation of
Scheme 1.2.A.2: Most prevalent methods of cyclic anhydride production.
31
epoxides (Scheme 1.2.A.2, reaction 6) has had more success with respect to
enantioselectivity, but it requires pre-synthesizing the chiral epoxide146 and in some cases
isolating the β-lactone intermediate.147,148 Specifically focusing on developing a direct and
inexpensive route to enantiopure versions of cyclic anhydrides through the use of chiral
ligands would therefore be highly advantageous, especially in their role as synthetic
precursors as current methods require special ordering from independent organic labs with
proportional prices and wait times. Such a method would enantioselectively synthesize
substituted cyclic anhydrides directly from olefins, CO2, and CO using a cheap and
abundant metal like nickel (ΔG ~ -6 kcal/mol if using ethylene), which would also generate
less waste and operate under milder conditions that are more tolerant of other functional
Scheme 1.2.A.3: (A) Achiral and (B) chiral nickel-promoted methods for the
production of cyclic anhydrides.
32
groups. The diversity of products available through functionalization of nickelalactones,
such as nickelamacrocyles, oligomers, carboxylic acids (including acrylic and propionic
acid), and muconic acid derivatives has been established.74,149-152 However, the nickel-
promoted production of cyclic anhydrides has been relatively unexplored, except for two
examples from coupling reactions77,153 and one using nickelalactones that were not CO2
derived (Scheme 1.2.A.3).154,155 There is currently no published work dealing with
enantioselective cyclic anhydride production from metal-promoted coupling of CO2,
alkenes, and CO. Based on the aforementioned precedent, our lab has developed several
nickel systems with the ability to couple CO2 with alkenes and/or alkynes and that
incorporate a variety of mono- or bi-dentate ligands, the most successful of which are of
the [N,N], [P,N], and [P,P] chelating types (Figure 1.2.A.1).114,115,153 These investigations
have determined a set of guidelines for ligand selection, namely that it must have an
electron-donating ability qualitatively equal to or greater than that of 2,2’-bipyridine to
effectively promote olefin binding and enough steric bulk to prevent significant formation
of the bis-ligand complex (e.g. [P,P]2Ni). The work of Greenburg et al. has indicated that
substituted cyclic anhydrides are feasible targets with such bidentate ligand scaffolds on
nickel,153 and the results presented here extend this concept to the use of chiral ligands
Figure 1.2.A.1: Examples of [N,N], [P,N], and [P,P] types of bidentate ligands. tBu =
tert-butyl, Cy = cyclohexyl.
33
(Scheme 1.2.A.3). Nine nickel complexes containing bidentate [P,P] chiral ligands were
screened for CO2/alkene/CO coupling ability, and it was found that two were competent
for MSA production, albeit with low enantioselectivities. Despite a disappointing lack of
selectivity, these investigations represent initial forays into steric ligand parameters that
could be key to coupling success and product control.
34
Figure 1.2.A.2: Structure key for ligand numbering. iPr = isopropyl, tBu = tert-butyl.
35
1.2.A.i.2: Results and Discussion
The nine chiral bidentate [P,P] ligands shown in Figure 1.2.A2 were selected based
on their electron-donating and sterically bulky groups, the coupling advantages of which
were previously discussed and where sterics can also be used to enforce regioselectivity.
They were combined with Ni(COD)2 as a nickel(0) source in order to form the
corresponding P2Ni(COD) complexes, and the majority of them were first screened for
CO2/alkene coupling ability using ethylene (Scheme 1.2.A.4, R’ = H) due to its simplicity,
as it had the least steric hindrance and avoided any complications due to multiple isomers.
Those complexes that were successful with respect to ethylene coupling were then tested
with propylene, a slightly more sterically hindered alkene (Scheme 1.2.A.4, R’ = CH3).
This added a substituent to the nickelalactone that allowed for both regio- and stereo-
isomers, however, analysis was simplified because subsequent CO insertion and product
elimination resulted in indistinguishable regioisomers. These coupling reactions were all
Scheme 1.2.A.4: Methods of nickelalactone production from (A) CO2/alkene coupling
and retrosynthetically from (B) a cyclic anhydride or (C) acrylic acid.
36
performed in J. Young (<4 atm combined gas pressure) or high-pressure (<13 atm
combined gas pressure) NMR tubes, and the optimized conditions for nickelalactone
formation are listed for each ligand in Tables A1.2 and A1.3 in Appendix 1.
Ligands 1 and 2 established the higher and lower steric limitations (respectively) of
the proposed reaction. Due to the size of its R groups, ligand 1 did not substantially bind
to nickel, forming very little of the 1-Ni(COD) complex (Figure A1.2) and decomposing
upon heating. Formation of the 2-Ni(COD) complex was successful, although a significant
amount of the bis-ligand species (2-Ni-2) was also formed due to the small size of the
methyl substitutents. It was found that 2-Ni(COD) did not bind significantly to ethylene,
even under harsh conditions (with heating or 11 atm of ethylene) (Figure A1.3). Attempts
to hydrogenate COD off of the nickel center using H2 alone had no effect, while addition
of H2 and ethylene only produced ethane and 2-Ni(COD). The methyl groups on 2 are
therefore not large enough to make the smaller ethylene molecule competitive with the
COD as a nickel ligand, which would force it off of the complex. As a result, no further
coupling attempts were made using ligands 1 or 2.
Although ligands 3, 4, and 5 formed the corresponding Ligand-Ni(COD)
complexes in good yields, initial CO2/ethylene coupling attempts were met with varying
degrees of success. Complex 3-Ni(COD) displayed very slow conversion to the 3-Ni-
ethylene complex and similarly slow, partial conversion to the nickelalactone upon
stepwise addition of ethylene and CO2 (Figure A1.15). These conversions were not
significantly enhanced by heating or higher partial pressures of ethylene and CO2. In the
case of 4-Ni(COD), there was similar partial lactone conversion upon CO2/ethylene
addition (Figure A1.16). Treatment with CO produced a small amount of SA on only one
37
occasion, indicating minimal coupling. The conversion from 5-Ni(COD) to the 5-Ni-
ethylene complex was kinetically fast but would have required unreasonably high gas
pressure to go to completion (Figure A1.17).
The behavior of these complexes suggested a thermodynamic barrier to coupling
that was overcome through the use of a Lewis acid, sodium tetrakis[3,5-
bis(trifluoromethyl)phenyl]borate (NaBAr4F). This solution was inspired by the previous
work of Jin et al., which indicated that the sodium cation stabilized the lactone through its
interaction with the two oxygen atoms, providing a thermodynamic sink that favored
lactone formation.114 While this was primarily applied as a method of inducing β-H
elimination in nickelalactone species, it was extended in this work to encourage CO2/alkene
coupling. Immediate improvement was observed upon addition of NaBAr4F to the coupling
reactions of 3-, 4-, and 5-Ni(COD) (Scheme 1.2.A.5). Using the sodium salt changed the
slow, partial conversion of 3-Ni(COD) to 3-Ni-lactone into full conversion after 24 hours
at 60 °C (Figure A1.4). In the case of 4-Ni(COD), which displayed almost no coupling
ability without NaBAr4F, nearly full conversion to 4-Ni-lactone was achieved after 12 hours
at 50 °C (Figure A1.6). The 5-Ni(COD) complex did not show as large of an improvement
as its conversion to 5-Ni-lactone was very low even in the presence of the Lewis acid
(Figure A1.8), however, this is still an advantage over the complete lack of coupling that it
showed beforehand. The identities of all three lactones were confirmed by the release of
SA upon CO addition (as seen by 1H NMR) and retrosynthetically by reaction of 3-, 4-, or
5-Ni(COD) with SA (Scheme 1.2.A.4, (B); Figures A1.18-A1.20). In the case of the 3-Ni-
lactone, the 31P{1H} NMR spectrum of the coupling product displayed major and minor
peaks, and the major product was identified as a γ-lactone using a combination of HSQC
38
and HMBC 2D NMR spectroscopy. The minor product was attributed to formation of the
β-lactone, but these peaks did not appear consistently and were occasionally observed pre-
CO2 addition, suggesting that they may have been impurities and not lactone-derived.
Further attempts to confirm via the reaction of 3-Ni(COD) with AA yielded large product
mixtures (Figure A1.21).
While these results are not sufficient to establish a full pattern with respect to steric
influences on the coupling ability of nickel complexes, the activity of 3-Ni(COD)
compared to 5-Ni(COD) (and to a lesser extent, 4-Ni(COD)) suggests that flexibility in the
ligand backbone may be advantageous to nickelalactone formation in these studies.
Therefore, ligand 6 was tested, as it has a similar two-carbon flexible chain and maintains
the general steric size and electron-donating ability of the phosphorous substituents, while
varying them from the five-membered phosphacycle ring featured in ligands 3-5. Ligands
7-9 introduced a ferrocene moiety (and consequently a larger bite angle) into the ligand
backbone, which in the cases of 7 and 8 represented a more rigid motif but with a variation
Scheme 1.2.A.5: Lewis acid enhancement of the 3-,4-, and 5-Ni(COD) coupling
reactions between ethylene and CO2.
γ-lactone β-lactone
39
in electronic effects, and in the case of 9 modified the linker distance and electronic
influence while maintaining the phosphacycle ring featured in ligands 3-5.
The 6-Ni(COD) species gave low conversion to 6-Ni-lactone (5 atm each gas, room
temperature, 48 hours), but this was attributed to an absence of sodium salt (Figure A1.10).
Product formation was confirmed up SA release after CO addition. Reactions involving
ligands 7-9 were all tested using the NaBAr4F Lewis acid additive. The 7-Ni(COD)
complex was formed in situ from Ni(COD)2 and 7 due to poor conversion at room
temperature and decomposition upon heating. While it showed evidence of binding
ethylene, it did not functionalize CO2 (even at 4 atm total gas pressure, Figure A1.12),
confirmed by HSQC and HMBC NMR spectroscopy using 13CO2. 9-Ni(COD) had a similar
problem in that it appeared to bind propylene (ethylene was not tested) but would not then
perform the oxidative coupling step with CO2 (Figure A1.14). 8-Ni(COD) could couple
ethylene and CO2 but only in very limited amounts (Figure A1.13), as evidenced by 13C
NMR spectroscopic analysis of the 13C-labeled SA product after CO release. Overall, the
nickel complexes containing ferrocene-based ligands exhibited a general lack of
CO2/alkene coupling ability and further trials with them were therefore not attempted. In
past works, nickelalactones containing these types of ligands were prepared via ligand
substitution from bipyridine-Ni-lactones rather than directly from coupling,86,114 so it is
possible that their capacity for electron-donation is insufficient to promote these types of
transformations.
Due to their CO2/ethylene coupling activities, complexes 3-, 4-, 5-, and 6-Ni(COD)
were employed in CO2/propylene coupling reactions (and all in the presence of Lewis
acid). Of these four compounds, neither 4- nor 5-Ni(COD) were able to form the
40
corresponding nickelalactones (4- and 5-Ni-lactone’, Figures A1.7 and A1.9). In the case
of 4-Ni(COD), this was due to the highly competitive formation of the CO2-adduct, to the
exclusion of the 4-Ni-lactone’ (even under low CO2 pressures). There was evidence of a
similar species being formed in the 5-Ni(COD) reaction, and 13CO2 studies indicated no
sign of CO2 functionalization or production of MSA upon treatment with CO. It is possible
that the rigidity of the ligand backbones in 4 and 5 prevented the necessary structural
rearrangements for incorporating a slightly more sterically hindered alkene into a ring with
CO2, instead favoring ligand displacement. In contrast, 3- and 6-Ni(COD) both
successfully coupled propylene, CO2, and CO to give MSA, as indicated by 1H and 13C
NMR spectroscopy as well as by chiral GCMS (Figures A1.5 and A1.11). The GCMS
analysis, however, indicated only a 1:1.3 ratio between enantiomers for 3 and a 1:2 ratio
(with slight peak overlap) for 6 (Figure 1.2.A.3). Therefore, while 3- and 6-Ni(COD)
presented superior coupling abilities compared to the other ligands tested, this did not
translate to any advantage with respect to enantiomeric selectivity.
Figure 1.2.A.3: Chiral nickel complexes competent for coupling propylene, CO2, and
CO to give MSA. iPr = isopropyl, tBu = tert-butyl.
41
1.2.A.i.3: Summary
There are numerous examples of nickel-promoted reactions of CO2 with olefins or
alkynes to give nickelalactones, and the incorporation of CO to release cyclic anhydrides
has been established. However, despite a demand for enantiomerically-pure cyclic
anhydrides, there has been no extension of the existing systems to include chiral species.
This work represents an initial foray into the enantioselective production of substituted
cyclic anhydrides using chiral nickel species for coupling alkenes, CO2, and CO. Of the
nine nickel complexes that were tested, only two (3- and 6-Ni(COD)) were competent for
MSA production from propylene, CO2, and CO, and then only with poor selectivities.
Nevertheless, this investigation uncovered a means of enhancing lactone formation through
the use of Lewis acid additives. In addition, these results suggested some important steric
ligand parameters, as well as a preference for a more flexible linker group between the
ligand binding sites. The greatest restriction to a more thorough examination of influential
ligand properties is the scarcity of commercially-available chiral ligands and the difficulties
associated with their synthesis, which prevents a systematic variation of ligand parameters
to determine what is needed for better enantiomeric control. The use of CO gas also
removes the option of catalysis (through the formation of catalytically inept L2Ni(CO)2),
which will be addressed in the following section.
42
1.2.A.ii.: The Use of CO Surrogates in Nickel-Promoted CO2-Alkene Coupling for the
Catalytic Production of Cyclic Anhydrides
1.2.A.ii.1: Introduction
The use of CO surrogates in cyclization and cycloaddition reactions to produce a
wide range of C=O functionalized species like ketones, esters, and aldehydes, as well as
heterocycles such as indolinones, lactones, and lactams has been established under both
metal-free156 and metal-catalyzed conditions.157,158,159 The incorporation of CO2 as a CO
surrogate has even been demonstrated, for example, in the palladium-catalyzed production
of aldehydes from aryl halides.160 However, the application of CO surrogates in the
catalytic generation of CO2-derived coupling products, such as cyclic anhydrides, has not
yet been explored. Instead, the few cases of using CO surrogates for making cyclic
anhydrides start from non-CO2 sources, such as carboxylic acids.161 The most closely
related example is the generation of α,β-unsaturated lactams, which was published by
Hoshimoto et al. in 2014 and used a (PCy3)2Ni(COD) catalyst (PCy3 =
tricyclohexylphosphine) for the [2+2+1] carbonylative cycloaddition of alkenes or alkynes
with imines (Scheme 1.2.A.6).162 Although this group had already established the
competency of this catalytic system for imine/alkyne coupling reactions,163-165 this is the
first nickel-catalyzed example of this carbonylative cycloaddition reaction using a CO
surrogate. The catalyst loading, however, was high (10 mol%; best TON achieved = 8) and
the imine substrates all contained a sulfonyl or tosyl moiety that would need to be
substituted out in an additional step to increase product diversity.162 Despite these
limitations, their coupling reaction is highly analogous to the nickel-promoted reaction of
CO2, alkenes, and CO previously discussed. Therefore, CO surrogates (regardless of
43
whether they are applied in a chiral or achiral setting) provide an opportunity to make an
otherwise stoichiometric reaction catalytic (Scheme 1.2.A.1) through the avoidance of
L2Ni(CO)2 (L = monodentate ligand) as a catalyst deactivation pathway. This section
presents experimental results concerning the viability of using CO surrogates to
catalytically liberate a substituted cyclic anhydride, MSA, from CO2/propylene coupling
using the nickel(0) species dcpe-Ni(COD) (dcpe = 1,2-bis(dicyclohexylphosphino)ethane).
This nickel complex was chosen because of its proven CO2 coupling ability.77,115 Overall,
the production of MSA was very low (a highest yield of 9%), even after testing three
separate CO surrogates with three different bases, with the major issue being the rate of
CO production compared to the CO2/propylene coupling rate.
Scheme 1.2.A.6: Previous application of CO surrogates with nickel species for the
production of cyclic lactams and the current proposed use for releasing cyclic
anhydrides.
44
1.2.A.ii.2: Results and Discussion
Originally, works incorporating CO surrogates like formates required directing
groups for the surrogate, ruthenium co-catalysts, strong bases, high temperatures and
pressures, large excesses of the surrogate, and often had limited substrate scopes and poor
yields.166-168 The groups of Manabe et al. and Tsuji et al. then discovered that weak organic
or inorganic bases could be used to evolve CO from CO surrogates in the absence of
transition metal reagents,169,170 greatly simplifying their application in catalysis. The CO
surrogates that were screened for the production of substituted cyclic anhydrides are listed
in Figure 1.2.A.4, along with the weak bases used with them in order to promote CO
release. Their selection is based on their multitude of appearances in previous
Figure 1.2.A.4: CO surrogates and bases used, with abbreviations.
45
works.157,158,162,169-172 Initially, stoichiometric reactions were attempted using multiple
equivalents of the CO surrogate/base pair on a small scale (a J. Young tube or scintillation
vial), with the intention of extending these studies to include catalytic reactions once
successful pairings had been established. Experiments began with phenyl formate (PF) and
triethylamine due to their successful application in Hoshimoto’s work. However, it was
found that CO release from PF was too kinetically fast for the propylene/CO2 coupling
reaction to compete, resulting in full and immediate conversion to the dcpe-Ni(CO2)2
complex without any formation of MSA (Table 1.2.A.1, entry 1). This issue was not
alleviated by the using of a different CO surrogate, N-formylsaccharin (NFS), or the use of
NFS even in the absence of a base promoter (Table A1.7). As a result, all further
investigations were performed starting with the pre-synthesized dcpe-Ni-lactone’ (the
nickelalactone from propylene/CO2 coupling). Reactions were initially still performed with
propylene and CO2 gas in the NMR tube headspace in order to allow for potential turnover.
The first attempt at generating MSA from dcpe-Ni-lactone’ and PF/triethylamine
yielded almost no product (Table 1.2.A.1, entry 2), with the primary problem still attributed
to the fast release of CO from PF. Efforts to control the CO release rate by varying the
polarity of the solvent using solvent mixtures were also unsuccessful (entries 3 and 4).
While the first reaction was performed in benzene (deuterated) due to its identity as the
optimized solvent for the previous work, further reactions in 30:70 and 50:50
benzene:tetrahydrofuran (THF) provided no enhancement, even with variations in CO
surrogate:base ([CO]:[B]) ratios (Table A1.8 and Figure A1.23). Performing the reactions
without any propylene or CO2 gas in case increased reaction pressures were inducing faster
CO release had little effect (Table A1.9). The greatest improvement for PF/triethylamine
46
Entry Ni
Species
CO
Surrogate
[CO]
Equiv.
[CO]
Base
[B]
Equiv.
[B]
Solvent MSA
Produced
(%)b
1 A PF 3 NEt3 4 THF-d8 0
2 B PF 3 NEt3 4 C6D6 0.05
3 B PF 3 NEt3 4 30:70
C6D6:THF-
d8
0.01
4 B PF 3 NEt3 4 50:50
C6D6:THF-
d8
0.03
5c B PF 9 NEt3 12 C6D6 0.09
6 B PF 9 NEt3 24 C6H6 2.00
7c B TCPF 4 NEt3 20 C6D6 9.00
8 B NFS 4 NEt3 10 C6D6 9.00
9 B PF 9 DBU 12 C6D6 -
10 B TCPF 3 DBU 4 C6D6 0.57
11 B NFS 3 DBU 4 C6D6 0.77
12 B PF 9 DMAP 12 C6H6 -
13 B TCPF 9 DMAP 12 C6H6 4.00
14 B NFS 9 DMAP 12 C6H6 6.30
15 B TCPF 4 - - C6H6 5.20
Table 1.2.A.1: Reaction of dcpe-Ni(COD) and dcpe-Ni-lactone’ Species with CO
Surrogatesa
aConditions: 6-10 mg Ni species, X equiv. [CO], and Y equiv. [B] in listed solvent in a J. Young NMR
tube with 0.5-1.5 atm each of propylene and CO2. Heat at 22-70 °C for 2 hrs-3 days. See Appendix 1 for
details. bMSA = methylsuccinic anhydride. Yield determined by GCMS using a naphthalene standard. cNo gas added.
47
was actually observed from increasing the [CO]:[B] ratio in 100% benzene solvent from
3:4 to 9:24 (Table 1.2.A.1, entry 2 versus 5 and 6), although this only brought the MSA
yield to 2%. Changing the identity of the CO surrogate gave the best results, with
NFS/triethylamine and trichlorophenyl formate (TCFP)/triethylamine both giving 9% yield
of MSA, albeit with different [CO]:[B] ratios. It was found that these ratios had to be
optimized based on the specific CO surrogate and base pairings (Tables A1.10-A1.12).
Changing the base paired with each surrogate gave no improvements, even though DBU
and DMAP were both tested in the nickel-catalyzed generation of cyclic lactams.162 In the
case of PF with either DBU or DMAP (Table 1.2.A.1, entries 9 and 12), there were
problems with product analysis due to overlapping GCMS peaks that could not be resolved.
For the combinations of TCPF (entries 10 and 13) or NFS (entries 11 and 14) with either
DBU or DMAP, the MSA yields only decreased. This was also the case with entry 15,
which repeated a reaction with a CO surrogate in the absence of base except with dcpe-Ni-
lactone’ as the starting nickel species. There was similarly no improvement, although it
confirms that these CO surrogates can produce CO even in the absence of base, making
control of the release rate difficult. Changing the reagent addition order such that the
surrogate was added last in order to minimize CO release before coupling was possible did
not have beneficial results. No evidence of MSA decomposition under the experimental
conditions (Figure A1.24), of MSA side reactions with byproducts from the CO surrogate
in the presence or absence of base (e.g. phenol, Figures A1.25 and A1.26), or of formation
of a deactivating nickel-CO2 adduct (Figure A1.27) was observed.
48
1.2.A.ii.3: Summary
Several combinations of organic CO surrogates and bases were used in an attempt
to mediate a CO slow-release mechanism to allow for catalysis of the nickel-promoted
production of substituted cyclic anhydrides from the oxidative coupling of alkenes, CO2,
and CO. Unfortunately, CO generation from the surrogates proved too fast for there to be
a competitive coupling reaction, even in the absence of base, and <10% of MSA was
produced from the dcpe-Ni-lactone’ starting material. One potential solution to this issue
would be to screen other nickel(0) species in order to find one that couples alkenes and
CO2 at a faster rate that could become competitive with the rate of CO production from the
surrogate. It has been established, however, that nickel complexes tend to exhibit greater
difficulty in the CO2/alkene coupling step compared to other metals (e.g. group VI), and
conditions like increased temperatures or gas pressures that might promote faster coupling
would also likely enhance the rate of CO release from a surrogate. Alternatively, ex situ
generation of CO in a simple two chamber system has been developed and could be applied
to control CO exposure.173 In this way, a separate surrogate solution could slowly bubble
CO into the lactone solution. This is a similar concept to that presented in a recent
publication by Hoshimoto et al., which indicated that they were able produce cyclic lactams
under very low pressures of CO gas (0.5 atm) in an autoclave without stirring using nickel
catalysts containing trisubstituted monodentate phosphines like tricyclohexylphosphine,
obtaining TONs of up to 100. They attributed this activity to the re-activation of their
L2Ni(CO)2 complex, but it could also be due to the reduced rate of diffusion of such a small
amount of CO gas in the larger volume of the autoclave.174 Regardless, these methods of
slowing down CO introduction even further are the most attractive options for finally
49
achieving nickel-based catalysis with respect to the production of cyclic anhydrides from
CO2.
50
1.2.B.: Nickel-Promoted Coupling of Isocyanates and Alkenes for the Production of N-
Substituted Acrylamides
1.2.B.1: Introduction
Isocyanates (R-N=C=O) are small, highly reactive organic molecules that are
isoelectronic with CO2.175,176 It is therefore feasible for them to participate in similar types
of reactions, except that these CO2 analogues have a much smaller thermodynamic barrier
to activation due to the substitution of a nitrogen (with five valence electrons) for an oxygen
(with six), and are much easier to functionalize as a result. Isocyanates are applied in a host
of organic transformations to give a wide assortment of products. For example, they contain
an electrophilic carbon which, when combined with nucleophiles, can yield carbonyl-
containing compounds (e.g. with alcohols or amines to give carbamates or ureas,
respectively) or cleavage products (e.g. with water to give amines and CO2).176 They are
also reactive intermediates, such as those formed by common rearrangement reactions like
Hofmann, Schmidt, Curtius, and Lossen.176,177 Isocyanate-derived products have a
correspondingly diverse range of applications, from the polyurethane industry178 to the
synthesis of diabetic drugs.179 Their metal-catalyzed reaction with olefins to give N-
substituted acrylamides in a manner parallel to that previously discussed for CO2, however,
is of particular interest.
Acrylamides are mainly used in the synthesis of polyacrylamides, which are water-
soluble thickening agents and flocculators that find applications in wastewater treatment,
paper production, ore processing, oil recovery, and the manufacture of permanent press
fabrics or dyes.180,181 Global demand for acrylamide (the simplest, unsubstituted
derivation) alone was 1.16 Mt per year in 2013,182 which does not even include the market
51
data for all of the specialty applications of substituted acrylamides, such as for temperature-
sensitive gels and resins.183-185 While this is not nearly the demand acrylic acid has, the
diversity of polymer products accessible through N-substituted acrylamides makes them
desirable targets.
Acrylamides are typically prepared from the reaction of acryloyl chloride with
amines186-188 or of acrylonitrile with water or aqueous sulfuric acid (Scheme 1.2.B.1).89,189
The first reaction does not require a catalyst, but the variety of substituents is limited and
produces chloride salts as waste.190 The second reaction with acid is catalyst-free but the
conditions are harsh, and with water requires either a reduced copper species189,191,192 or
immobilized nitrile hydratase to catalyze the process.193 The former gives poorer yields,
deactivates quickly, and can also catalyze the production of unwanted side products like
polymers or acrylic acid,193-196 so the latter is preferred. In both cases, the starting materials
require multi-step pre-synthesis and isolation, where the acryloyl chloride is derived from
acrylic acid and a chloride source and the acrylonitrile is from propylene, ammonia, and
oxygen. Developing a more direct method from a C1 carbon source to the intended
Scheme 1.2.B.1: Large-scale methods of synthesizing acrylamides.
52
acrylamides, such as from isocyanates and alkenes, would therefore be economically
beneficial and produce less waste.
The inherent activity of isocyanates permits a multitude of metal-free
transformations with strong nucleophiles, however, their interactions with alkenes and
alkynes are metal-promoted (with the exception of a very few, specific, highly reactive
cases that do not produce acrylamides).197,198 While there are many examples of metal-
catalyzed coupling reactions between isocyanates and unsaturated compounds like
aldehydes,199-201 imines,199, and alkynes202-204 that feature promotors incorporating a variety
of metals (such as osmium, rhodium, iridium, platinum chromium, tungsten, copper, iron,
and nickel),205 those between isocyanates and alkenes are dominated by nickel-containing
complexes.205,206 The earliest and most pivotal work in this area was done by Hoberg
beginning in the 1980’s, in parallel to his studies of CO2/alkene coupling on nickel. The
first publications of Hoberg et al. on this topic established the successful formation of the
corresponding azanickelacyclopentanone from ethylene and phenyl isocyanate followed
by hydrolysis to give propanilide or CO insertion to give N-phenyl succinimide.207,208 They
built upon these preliminary results and were able to release the desired N-phenyl
acrylamide by heating (80 °C), but at this stage the reaction only proceeded
stoichiometrically and had to be performed at very low temperatures (-78 °C).209 Soon
after, they were able to achieve catalysis by using slightly higher ethylene pressures and a
more electron-donating ligand on nickel, but their TONs were very low (<20) and they had
a mixture of acrylamide products from both single and double insertion of ethylene into
the metallocycle.210 The scope of the reaction was extended to include a variety of alkenes,
including linear alkenes of the type R1R2C=CH2 (where R1 = Ph, OEt, SPh, CO2Me, CH3,
53
and CF3 when R2 = H or R1 = R2 = F),211-214 cyclic alkenes (e.g. cyclopentene),215 dienes
(e.g. 1,3-butadiene),216-218 and bicyclics,219 but in all cases the other reactant was phenyl
isocyanate or its derivatives (i.e. one example also used 4-methylphenyl isocyanate). There
was also no significant increase in TONs. The incorporation of additives to divert the
reaction to other products was also investigated, which includes the aforementioned use of
acid to give a saturated amide207,212,213 or CO to yield a succinimide,207,211,212 as well as the
use of oxidizing agents (e.g. FeCl3 or I2) to produce α,ω-diamides.220
The proposed reaction sequence for nickel-catalyzed isocyanate/alkene coupling is
shown in Scheme 1.2.B.2, where (under the right experimental conditions) acrylamide
release occurs through thermally-induced β-H elimination from the nickel metallocycle.209
This represents a drastic change from nickelalactone behavior, suggesting that replacing a
Ni-O bond with a Ni-N bond allows for greater structural flexibility in the metal ring. In
this way, acrylamide formation via nickel-induced coupling reactions bypasses the issues
of CO2/alkene coupling while still accessing valuable functionalized products.
For a substituted alkene like propylene, two potential coupling products could be
formed based on the position of the alkene R group on the nickelacycle (α- or β- to the
metal center, Scheme 1.2.B.2). Location at the α-carbon gives disubstituted α,β-unsaturated
amides upon β-H elimination, while being at the β-position yields a 1,1-disubstituted
acrylamide.206 In Hoberg’s studies, which incorporated simple trisubstituted alkyl or aryl
phosphine ligands like tricyclohexylphosphine or triphenylphosphine, the major species
contained the R group at the α-position,206,211,212 except in the case of fluoroalkenes.213
Later work by Schleicher and Jamison demonstrated that by using much bigger, sterically
bulky ligands like N-heterocyclic carbenes (NHCs) on nickel (the coupling ability of which
54
had already been established for isocyanates and alkynes),221 the regioselectivity could be
forced to favor the position of the R group at the β-carbon to exclusively give 1,1-
disubstituted acrylamides. While this is an impressive example of regio-control and they
achieved good yields (up to 86%) and selectivities (up to 100%), their catalyst loading was
high (10 mol%), so their best TON was 9.190
With the exception of the research presented above by Jamison’s group, there has
been very little advancement in the field of nickel-catalyzed isocyanate/alkene coupling
reactions for acrylamide production since Hoberg’s initial work, especially in terms of
increasing catalytic activity.202,205 Newer studies have instead focused on nickel-catalyzed
Scheme 1.2.B.2: Nickel-catalyzed isocyanate/alkene coupling reactions with
conditions for accessing specific acrylamide regioisomers as well as propionamides and
succinimides.
55
cycloaddition reactions between isocyanates and unsaturated compounds (alkynes, allenes,
dienes, acrylates) to give heterocycles, in generally low yields (TONs <20) and often with
poor selectivities (due to potentially multiple isocyanate insertions).222-229 This section
concerns the application of a well-established nickel complex for CO2/alkene coupling
(dcpe-Ni(COD); dcpe = 1,2-bis(dicyclohexylphosphino)ethane), toward the coupling of
ethylene with ethyl, phenyl, and tert-butyl isocyanate to give N-substituted acrylamides
(Scheme 1.2.B.3). By varying steric and electronic factors on the isocyanate substituents,
it was hoped that a better understanding of reaction promoting parameters would be
obtained. Overall, it was found that while dcpe-Ni(COD) successfully performed the
coupling reactions, the greatest obstacle to catalysis was the highly competitive isocyanate
trimerization pathway to form isocyanurates. This resulted in only trace amounts of the
intended acrylamide or succinimide species and which ultimately could not be overcome.
1.2.B.2: Results and Discussion
Initial experiments were focused on evaluating the coupling ability and product
selectivity of the nickel species, so they were performed as stoichiometric reactions in
either J. Young (total gas pressures <4 atm) or high-pressure (total gas pressures <13 atm)
NMR tubes. All isocyanates were first tested with ethylene because it is sterically
undemanding and avoids the complication of introducing regioisomers. These un-
Scheme 1.2.B.3: Nickel-catalyzed isocyanate/alkene coupling reactions pursued in this
work.
56
optimized conditions, however, do not preclude the extension of any preliminary results to
encompass catalytic reactions using substituted alkenes after coupling competency had
been established. Even though the coupling activity of isocyanates usually increases with
the electron-withdrawing nature of the group on the nitrogen (due to the displacement of
electron density onto the R group upon nucleophilic attack), the isocyanate screening began
with ethyl isocyanate because of its simplicity and moderate steric bulk. It also has a low
boiling point, which allows for its addition to the reaction solution via vacuum transfer
after ethylene pressurization. As nickel-ethylene complexes are often unstable in the
absence of a headspace of ethylene, this process ensured the formation of the nickel-
ethylene species prior to isocyanate introduction, which could contribute to favoring the
formation of the desired coupling product.
The stepwise addition of 2 atm of ethylene and 1 equivalent of ethyl isocyanate to
10 mg of dcpe-Ni(COD) began with the promising appearance of a new set of doublets in
the 31P{1H} NMR spectrum (Figure A1.28). These would be consistent with the
inequivalent phosphine groups found in the five-membered nickel metallocycle formed
from ethylene/isocyanate coupling. However, the reaction did not proceed to completion
at room temperature, and upon heating to 60 °C and then 70 °C, multiple phosphorous-
containing products began to appear (Figure A1.28). Adding CO at the end to release any
nickel-bound organic compounds resulted in a mixture of products, some of which
contained olefinic protons (Figure A1.29). Suspecting that 1 equivalent of isocyanate was
perhaps insufficient for the reaction to proceed fully, the experiment was repeated in a
similar manner using 10 equivalents of ethyl isocyanate. After several days at room
temperature, the same (multiple) nickel species had returned (even in the absence of heat;
57
Figure A1.30), and the solvent was removed under vacuum. The resulting mixture was
recrystallized from pentane/toluene to yield a solid and a yellow solution. Once the yellow
solution was dried under vacuum, the two solids were independently dissolved in C6D6 and
treated with CO gas. The solid from the pentane/toluene recrystallization was N-ethyl
succinimide, while that from the supernatant was N-ethyl acrylamide, as confirmed by
comparison of the 1H NMR spectra (Figures A1.31 and A1.32) to authentic compounds
and by GCMS analysis. Reaction of dcpe-Ni(COD) with purchased N-ethyl acrylamide
yielded a nickel complex with a set of two phosphorous peaks that were very close in
chemical shift to one of the coupling products, tentatively assigned as the η2-acrylamide
species (Figure A1.33). These findings, however, only accounted for two of peak sets in
the 31P NMR spectrum.
In order to verify that all of the nickel species appearing in the NMR spectrum
contained ethylene, 13C-labeled ethylene was used in a coupling experiment under similar
conditions (~1 atm 13C-ethylene, 10 equivalents of ethyl isocyanate). It was immediately
observed that only two of the four sets of phosphorous peaks indicated incorporation of the
13C-ethylene through extra splitting due to 13C-31P interactions (Figure A1.34). Through a
combination of 2D (HSQC, HMBC, COSY) and 1D (1H, 31P, 13C) NMR experiments, it
was determined that the other two phosphorous-containing reaction species (accounting for
the major reaction products) did not incorporate any ethylene (Figure A1.35). Assignments
for the metallocycle and η2-acrylamide nickel complexes formed by ethylene/ethyl
isocyanate coupling are shown in Figure 1.2.B.1. The other two species were attributed to
58
isocyanate-associated nickel species on the pathway to isocyanate dimerization and/or
trimerization, which are well-known reaction behaviors of isocyanates.176,230-235 Peaks
potentially belonging to the isocyanurate cyclotrimerization product were also identified
in the 1H NMR spectrum (Figure A1.36).
Attempts to minimize or suppress the isocyanate cyclotrimerization pathway
included increasing the ethylene:isocyanate ratio (both in an NMR tube and using a Parr
reactor), performing sequential additions of isocyanate as a method of decreasing the
average isocyanate concentration, and using base additives to try to promote the production
Figure 1.2.B.1: 31P{1H} NMR spectrum in C6D6 of the nickel-promoted coupling
products from the ethylene/ethyl isocyanate reaction (2 atm/10 equiv.). Products C/D
do not incorporate ethylene. 300 MHz spectrometer.
E E F F
C D C D
A
B
59
of free acrylamide. Using a combination of 8 atm ethylene with 2 equivalents of ethyl
isocyanate in a high-pressure NMR tube reduced the observed products to the nickel η2-
acrylamide and one of the nickel-isocyanate adducts in an approximately 50:50 mixture
(Figure A1.37). As this was nearing the pressure limitations of the NMR tube, further
reactions were performed in a Parr reactor (Table A1.14). It was found that pressures of
~1,000 psi ethylene successfully stopped the cyclotrimerization process, however, the
reactions yielded no acrylamide or any other recognizable nickel species along the coupling
pathway (entries 9-11, Figures A1.38 and A1.39). A possible reason for this would be
multiple ethylene insertions into the five-membered nickelacycle. At pressures above 600
psi ethylene, no acrylamide or acrylamide-containing nickel species were observed (entries
6-11), but at pressures between 300-600 psi ethylene, the cyclotrimerization pathway still
remained active (entries 1, 3-5, Figures A1.38 and A1.39) unless the equivalents of
isocyanate were so low that no reaction occurred (entries 2 and 9). Adjustments in the
number of isocyanate equivalents yielded no significant improvements in selectivity
(entries 2 vs. 3, 6 vs. 7, and 10 vs. 11). Suspecting that product inhibition might be
contributing to the generation of the isocyanurate as the major organic species, a sodium
phenoxide base additive (sodium 3-fluorophenoxide) was used to try to encourage
acrylamide release from the nickel η2-acrylamide complex. Unfortunately, this did not
result in greater acrylamide yields (Table A1.13, entry 4). Attempts to induce free
acrylamide from the nickel η2-acrylamide bound species using other types of bases (e.g.
sodium tert-butoxide, 1,8-diazabicyclo[5.4.0]undec-7-ene = DBU) were also unsuccessful.
One coupling attempt was made using 6-Ni(COD) (see Figure 1.2.B.4 for structure) to
60
ascertain if a change in ligand would enhance selectivity, but this yielded no improvements.
Further experiments with ethyl isocyanate were therefore not pursued.
In an effort to discourage the cyclotrimerization reaction, tert-butyl isocyanate was
employed in the ethylene coupling reaction due to its sterically more hindered R group.
The first coupling experiment gave two major phosphorous-containing species in the
31P{1H} NMR spectrum after gentle heating (60 °C) over several days (Figure 1.2.B.2),
without any sign of an isocyanurate in the 1H NMR spectrum. Addition of CO gas gave the
free N-(tert-butyl) acrylamide in very small yield, the identity of which was verified by
both 1H NMR spectroscopy (Figure A1.40) and GCMS analysis. Confirmation of the
concurrent production of N-(tert-butyl) succinimide was obstructed by the NMR peaks
corresponding to the ligand. However, the position of one of the phosphorous peak sets is
consistent with the formation of the tert-butyl isocyanate/ethylene derived nickel
metallocycle if compared to the analogous species prepared from ethyl isocyanate and
Figure 1.2.B.2: 31P{1H} NMR spectrum in C6D6 of the nickel-promoted coupling
products from the ethylene/tert-butyl isocyanate reaction (2 atm/10 equiv.) See Figure
1.2.B.1 for labeling guide. 300 MHz spectrometer.
F F E E
B
61
ethylene. Despite this promising start, future coupling experiments using tert-butyl
isocyanate were unfortunately unable to reproduce these results outside of one occasion, in
which the yield of acrylamide was similarly limited. In all other cases, coupling did not
occur and the reaction mixture sat at the dcpe-Ni-ethylene complex, even with low ethylene
pressures and high isocyanate equivalencies at elevated temperatures. An experiment
featuring stepwise reagent addition beginning with tert-butyl isocyanate indicated the
formation of a nickel-bound isocyanate species. However, the isocyanate ligand was
immediately displaced upon ethylene addition (Figure A1.41). If the mechanism for the
coupling pathway requires initial binding of the isocyanate before it can proceed, the ease
with which ethylene displaces tert-butyl isocyanate may inhibit the intended reaction. In
addition, rapid decomposition and polymerization of tert-butyl isocyanate was observed
under preparative and storage conditions, so it was hypothesized that this instability may
account for its limited capacity to engage in coupling interactions with ethylene to a
significant degree.
The final reactive ethylene coupling partner to be applied to this transformation was
phenyl isocyanate, with which Hoberg found great success in both the production of
nickelacycles and the thermally-induced elimination of N-phenyl acrylamide. Initial
studies indicated full conversion from the dcpe-Ni-ethylene complex to one major nickel
species (and giving only a trace amount of a second) after only 1 hour at room temperature
with only 1 equivalent of phenyl isocyanate (Figure 1.2.B.3). The 31P spectral peaks of this
new compound appeared as singlets, which is consistent with those of the nickelacycles
formed from the coupling reactions between ethylene and the other two isocyanates.
Conversely, the nickel η2-acrylamide species have all displayed large enough coupling
62
constants for their phosphorous NMR peaks to appear as two distinct doublets (and which
likely accounts for the identity of the other trace complex appearing in the NMR spectrum).
No clear vinylic peaks appeared in the 1H NMR spectrum either before or after addition of
CO gas, and no acrylamide was detected by GCMS. However, these analysis methods did
indicate the presence of N-phenyl succinimide (Figure A1.42). The results suggested that
while the coupling reaction between ethylene and phenyl isocyanate was comparatively
facile, elimination of the acrylamide was more difficult than for the other isocyanates. This
barrier could not be overcome by simple heating, as was the case with Hoberg’s
experiments.209
In a manner analogous to the previously mentioned methods aimed at inducing β-
H elimination in nickelalactones for the formation of acrylates, a Lewis acid and several
bases of varying strengths were employed to try to force acrylamide generation from the
nickelacycle. In the case of running the coupling reaction in the presence of a Lewis acid,
Figure 1.2.B.3: 31P{1H} NMR spectrum in C6D6 of the nickel-promoted coupling
products from the ethylene/phenyl isocyanate reaction (1 atm/2 equiv.). See Figure
1.2.B.1 for labeling guide. 300 MHz spectrometer.
F F
63
NaBAr4F, there was no evidence for production of the nickel η2-acrylamide species or free
acrylamide by NMR spectroscopy or GCMS (Figures A1.43 and A1.44). It therefore
appeared as though this Lewis acid was not capable of significantly disrupting the Ni-N
bond in the nickelacycle to allow for ligand re-orientation and subsequent β-H elimination.
The addition of triethylamine as a weak base similarly had no effect (Figure A1.45).
Reasoning that perhaps triethylamine was simply not basic enough to deprotonate the β-
hydrogen of the nickelacycle, a slightly stronger base was utilized. The reaction revealed
some promising olefinic peaks when performed in the presence of DBU, (Figures A1.46
and A1.47), however, a 2D HSQC NMR spectrum indicated that there were no vinylic
methylene peaks in the product generated after treatment with CO, and no acrylamide was
identified by GCMS. This would suggest that while DBU induces a breakdown of the
nickelacyle, it also causes further reaction either at the nickel center or of the desired
product. Adding hydrochloric acid gas to the compound generated by DBU in order to
assess its reactivity gave no other insight into its identity, although its associated olefinic
peaks did disappear from the NMR spectrum. While this verifies the presence of an
unsaturated species, it was not the targeted acrylamide, so the effect of using progressively
stronger bases was then tested. It was found that sodium 3-fluorophenoxide was unable to
initiate the desired elimination (Figures A1.48 and A1.49). With NaOtBu as an additive,
however, a new set of olefinic peaks appeared in the proton spectrum, potentially as part
of the vinylic set of N-phenyl acrylamide. A combination of HSQC and COSY NMR data
obtained using the reaction solution (Figure A1.50) indicated that these peaks correlated to
a –CH=CH2 functional group consistent with an acrylamide. Encouraged by these results,
enhancement of the coupling reaction and hence the acrylamide yield was attempted
64
through the use of a Parr reactor in order to accommodate higher ethylene pressures and
equivalents of phenyl isocyanate. Under these conditions (300-500 psi ethylene, 20-50
equivalents phenyl isocyanate, 20 equivalents NaOtBu, 60 °C, 16 hours), no apparent
acrylamide was generated. Consequent NMR-scale experiments revealed the reaction of
phenyl isocyanate with NaOtBu, the rate of which is likely accelerated when combined
with the excesses of base and increased pressure conditions of the Parr environment.
Therefore, while NaOtBu is a competent base for inducing β-H elimination and product
release from more reticent nickelacycles derived from alkenes and isocyanates, it is
incompatible with the conditions necessary for exploring catalysis and would only be
suitable for generating acrylamides on sub-stoichiometric scales.
Due to Hoberg’s achievement of β-H elimination from an ethylene/phenyl
isocyanate-derived nickelacycle,209 the final attempt to devise an alternative way of forcing
off the acrylamide focused on the use of other ligands (depicted in Figure 1.2.B.4.). Four
Figure 1.2.B.4: Alternative ligands tested in Ligand-Ni(COD)-promoted
ethylene/isocyanate coupling reactions.
65
other Ligand-Ni(COD) complexes were tested (including 6-Ni(COD) from section
1.2.A.i.) for coupling ability with respect to ethylene and phenyl isocyanate and their
enforcement of β-H elimination from the metallocycle intermediate. They were chosen
based on their previously-demonstrated CO2/alkene coupling capabilities, and in the case
of BenzP*-Ni(COD), catalytic activity. For this application, however, none of these
complexes yielded acrylamide, instead giving results that ranged from unreactive (BenzP*)
to partial activity (6, giving the metallocycle but no acrylamide) and decomposition (MAP,
P2N2). A more systematic ligand variation would likely be required to fully understand the
factors promoting or inhibiting the sought-after elimination.
1.2.B.3: Summary
A series of experiments aimed at leveraging the inherent reactivity of isocyanates
in nickel-promoted alkene transformations was undertaken, with the ultimate goal of
significantly enhancing the catalytic production of N-substituted acrylamides in this
respect. It was instead found that diversifying the pool of isocyanate starting materials
elicited a range of challenges, which perhaps influence the prevalence of certain
isocyanates in alkene coupling reactions in the literature (e.g. phenyl isocyanate). In the
case of isocyanates possessing a small, electron-donating R group like ethyl, a
cyclotrimerization pathway that was highly competitive with the intended coupling
transformation was activated, and conditions that suppressed this side reaction were
similarly detrimental to acrylamide formation. Employing a larger R group, like tert-butyl,
effectively discouraged cyclotrimerization but then provided too great an obstacle to
coupling. Phenyl isocyanate was not inclined to cyclotrimerize and it easily underwent an
oxidative coupling reaction with ethylene at a nickel center, fully converting to the
66
metallocycle in short time periods under mild conditions. However, it was then very
difficult to induce product elimination, even with strong base additives that are also
unsuitable for catalysis. Thus each isocyanate faced a unique challenge at a different stage
in the coupling reaction with alkenes. Future work should approach these problems from a
ligand perspective. By instead starting with a nickel complex that is known for spontaneous
or thermally-induced acrylamide release (where the ligand is reminiscent of those
established by previous work, such as NHC’s or a PR3 group) and then assessing any issues
that occur through variation of the identity of the isocyanate, factors that control the
intended behavior can be more easily identified and controlled. Alternatively, a practical
experimental solution could involve enacting a slow addition method for introducing the
isocyanate into the reaction mixture continuously but gradually over time. This may help
discourage polymerization pathways while still benefitting from the ease of coupling and
product elimination advantages that were displayed by ethyl isocyanate.
1.2.C.: Section 1.2 Summary
Nickel has an illustrious history with respect to coupling CO2 (and CO2 analogues)
with olefins and alkynes, with the consequent production of a vast array of commercially
valuable chemicals from an inexpensive, widely available carbon feedstock. The use of a
cheaper, more abundant metal like nickel also represents an economically beneficial
transition from precious metals like palladium or ruthenium. The work presented herein
sought to extend previously established methods for the transformation of CO2 to include
specialty targets like enantiopure substituted cyclic anhydrides and N-substituted
acrylamides. While ultimately these goals were not achieved, many of the limiting factors
were identified, which provides a basis for developing ways to overcome these obstacles.
67
The outlook for devising and improving upon the nickel-catalyzed production of these
compounds is therefore promising.
68
1.3: Transition-Metal-Catalyzed Hydrogenation of CO2 to Methanol
1.3.1: Introduction
In addition to the previously discussed metal-catalyzed coupling reactions between
CO2 and unsaturated organic molecules for the generation of new C-C bonds, a second
means of CO2 activation and transformation is through its reduction using dihydrogen (H2).
Depending on how many equivalents of H2 are incorporated, CO2 hydrogenations can yield
a variety of C1 products, including CO, formic acid, methanol, and methane (Scheme
1.3.1).236,237 Two that have been of particular interest as target molecules are formic acid
and methanol, due to the market for the chemicals themselves, their potential as hydrogen
storage materials, and the feasibility of their reaction thermodynamics236,238-240,266 The
following section will focus on the transition-metal-catalyzed production of methanol from
CO2 and H2.
The global market demand for methanol has seen significant recent growth (49 Mt
in 2010 to 80 Mt in 2016)27,28,241, which has been attributed to its incorporation into newer
end-uses (such as its conversion to olefins for the production of plastics) and the increasing
Scheme 1.3.1: Products of CO2 hydrogenation reaction and the Gibbs free energy for
formic acid and methanol generation.
ΔG° = -2.4 kcal/mole
ΔG° = 7.8 kcal/mole
69
popularity of methanol or methanol-derived compounds (like dimethyl ether) for fuel
blending or fuel substitutes as a way of extending gas resources.238,241 The methanol-to-
olefin process alone has gone from accounting for only a trace amount of methanol
production to being the second largest application of methanol by end-use (an 18% increase
over a 5 year timeframe).27,241 This trend reflects the fact that despite its widespread
recognition as an organic solvent, >95% of the methanol produced annually is actually
converted into other valuable chemicals.27 Generating methanol from CO2 can therefore be
viewed as a means of accessing a host of other functionalized species from a carbon source
that might otherwise be unavailable to them by direct reaction (see Chapter 2 for a more
detailed discussion of this topic). For example, formaldehyde and acetic acid account for
over 35% of the global methanol demand (formaldehyde far exceeding acetic acid at 27%
and 9%, respectively),27 and they are subsequently converted into other high-demand
products like paints, adhesives, textiles, films, resins, plastics, and pharmaceuticals.242-244
The economic benefits of sourcing methanol from an inexpensive, non-toxic, and abundant
carbon feedstock like CO2 would therefore be realized in a diverse array of industries.
As mentioned briefly in the introduction to this chapter, industrial-scale methanol
production currently relies on the use of an alumina-supported mixed copper/zinc oxide
Scheme 1.3.2: Industrial method of methanol production from syngas, depicting (1) the
component equilibria and (2) the overall reaction.
70
heterogeneous catalyst for the conversion of syngas (a mixture of CO, CO2, and H2)
through the series of reactions shown in Scheme 1.3.2.245,246 Unfortunately, the co-
production of water necessitates a time and energy-intensive post-synthetic distillation
process for the isolation of the desired methanol.247 While highly efficient, the harsh
temperature and pressure conditions of the methanol production process are incompatible
with many other functional groups, prohibiting the use of a one-pot method for
streamlining further transformations. Accessing methanol using metal-catalyzed CO2
hydrogenation under milder conditions would make this goal much more feasible.239,240
Additionally, syngas is mostly derived from non-renewable coal and natural gas
resources,242 with only a small percentage available via low-yield248,249 biomass
fermentation.250-252 The incorporation of CO2 as a methanol C1 carbon source therefore
represents the opportunity for a much-needed transition from fossil-fuel reliance to a large-
scale renewable resource, imparting environmental as well as economic benefits.
Developing this methanol production method to integrate H2 derived from solar,
geothermal, or hydro powered electrolysis would maximize its sustainability, a process
which has in fact been recently realized on an industrial scale with the opening of the first
renewable CO2-to-methanol plant in Iceland in 2012. Its production capacity in 2015 was
over 5 million liters per year and it recycles an estimated 5.5 thousand metric tons of CO2
annually.253
In addition to its role as an intermediate on the pathway to other commercial
chemicals, methanol is an attractive potential hydrogen storage material. Methanol is 12.6
weight percent hydrogen, making its hydrogen storage capacity about three times that of
the other popular CO2 hydrogenation product, formic acid (4.4 weight percent).239 A major
71
challenge to the practical application of fuel cells as alternative energy sources is the
necessity of transporting a highly explosive compressed gas, which is both dangerous and
incompatible with the current infrastructure designed for storing, transporting, and
distributing liquid fuels.238 By employing a catalyst for CO2 hydrogenation to methanol,
H2 can be safely stored as a stable, easy to transport liquid until it is needed, at which time
methanol dehydrogenation is initiated to release the H2 (and CO2 enters back into the
storage cycle; Scheme 1.3.3). This has the added benefit of providing energy with no net
CO2 production. The concept of storing and transporting energy in the form of chemical
bonds through a reversible cycle of bond-making and bond-breaking is also relevant to
alternative energy sources like solar or wind, which produce power intermittently. In order
to meet a continuous demand, the excess electrical energy generated by these processes at
a given time could be used in chemical synthesis, thereby storing it as chemical energy to
be released when their production levels are low or nonexistent.238
Extensive studies into metal-free and metal-catalyzed CO2 hydrogenation to
methanol have been reported.239,240 Despite the use of frustrated Lewis pairs237,254,255 and
N-heterocyclic carbenes256,257 as catalysts for this transformation, the majority of studies
concern the use of transition metal catalysts due to generally superior activities.238,239,240
Scheme 1.3.3: General method of H2 storage using reversible metal-catalyzed CO2
hydrogenation/dehydrogenation.
H2
CO2
HCO2H
CH3OH
Hydrogenation Dehydrogenation H2
Potentially
renewable
source
Fuel Cell
Commodity chemicals
[M] [M]
72
While a multitude of reductants (with or without metal catalyst) have been successfully
applied to this reaction (including boranes,258,259 silanes,257,260 and 1,2-dihydropyridine261),
H2 represents the less expensive option with the least waste generation.239,262 Finally,
heterogeneous CO2 hydrogenations require harsh conditions (>250 °C) and often also
display low activities and poor product selectivity (such as partial CO2 reduction to yield
CO, which is a poison for fuel cells).240,263 Therefore, only homogeneous transition-metal
catalysts for CO2 reduction to methanol using H2 will be presented.
The overall conversion from CO2 to methanol is (slightly) thermodynamically
favorable, with a ΔG = -2.4 kcal/mole.236 However, there is a significant kinetic activation
barrier to the direct hydrogenation of CO2 to methanol that would require the use of very
harsh conditions (e.g. highly elevated temperatures and pressures) to overcome, even in the
presence of a catalyst. These conditions also provide access to undesirable side products
like CO or methane.239 For example, the first case of CO2 hydrogenation to methanol was
reported in 1993 by Tominaga et al. using a [Ru3(CO)12] catalyst with KI at 240 °C and 80
bar (1200 psi) total pressure to give a TON of only 95. The selectivity was also very low,
producing a mixture of CO, methane, and ethane in addition to methanol.264 Future efforts
using other metal carbonyl catalysts (containing M = Ir, Rh, W, Mo, Co, and Fe) did not
overcome these limitations, with Ru providing the best results.265 Consequently, most work
in the field of CO2 hydrogenations takes advantage of CO2 derivatives (e.g. formic acid,
carbonates, esters, amides, carbamates, and ureas) with lower activation barriers to
hydrogenation in order to carry the reaction forward stepwise under much milder
conditions and with greater selectivities (Scheme 1.3.4).236,239 Many of these compounds
are formed upon CO2 sequestration,240 so their incorporation into CO2 hydrogenation
73
pathways provides an opportunity for a smooth transition from CO2 capture to
transformation and utilization. Hydrogenation of CO2 to formic acid is thermodynamically
uphill by ~7 kcal/mole,266 which can be challenging for many systems. Consequently,
nearly all examples of generating methanol by this pathway do so starting from formic acid
rather than CO2,239,267,268 with the exception of only one recently developed iridium
catalyst.269 There are similar hydrogenation examples of generating methanol from
formate, carbonate, carbamate, and urea starting materials270-272 (covered in detail in the
cited reviews).239,240 While these cases establish proof of concept in the feasibility of using
CO2 derivatives to facilitate CO2 hydrogenation reactions, the necessity of independently
Scheme 1.3.4: Methods of indirect hydrogenation of CO2 to methanol.
74
synthesizing and isolating each species detracts from the appeal of this method on the basis
of cost and efficiency. In addition, the challenges with transitioning a two-pot method (CO2
to derivative, derivative to methanol) into a one-pot sequence are non-trivial. It is instead
preferable to leverage the kinetic advantages of these derivatives in situ by using additives
(e.g. amines or alcohols) to generate the CO2 derivative as an intermediate in a one-pot
reaction. With only a few exceptions, previous applications of this strategy of shuttling
CO2 to methanol can be divided based on whether they proceed via an ester or a formamide
intermediate (Scheme 1.3.5).
Independent investigations by the groups of Sanford and Leitner established the
field for this manner of achieving a more level kinetic profile for CO2 hydrogenation, which
encompasses a rather limited number of overall examples. Sanford et al. first presented a
cascade catalysis method for achieving CO2 hydrogenation to methanol employing three
different catalysts in a one-pot, three-step transformation (Figure 1.3.1, (1)).262 CO2 was
first hydrogenated to formic acid using ruthenium catalyst A, followed by conversion to
methyl formate (an ester) using scandium catalyst B, and finally hydrogenation of the ester
to methanol using ruthenium catalyst C. Generation of the intermediate ester in these types
Scheme 1.3.5: Metal-catalyzed CO2 hydrogenation to methanol via (1) esters or (2)
formamides.
75
of reactions is achieved by the addition of an alcohol, often methanol, which should be re-
generated at the end along with a new molecule of methanol. Initial activity (TONs < 3)
was limited by the deactivation of catalyst C due to high CO2 pressures and by the scandium
triflate catalyst for step two. By devising a two-vessel setup where the first two steps
occurred in the inner vessel, followed by volatilization of the ester to an outer vessel
containing catalyst C in solution, they were able to achieve TONs of up to 21 for the
selective production of methanol at 75-135 °C and 40 bar (600 psi) total gas pressure.262
Overall, these TONs were quite low, and using three incompatible catalysts in a two-vessel
setup would be highly impractical. However, these results represented a significant step
forward in terms of increasing product selectivity and decreasing temperature and pressure
requirements.
One year later, Klankermayer and Leitner reported the one-pot hydrogenation of
CO2 to methanol (selectively), achieving TONs of up to 221 using a single ruthenium-
triphos catalyst (Figure 1.3.1, (2)) with a co-catalytic amount of an acid
(bis(trifluoromethane)sulfonamide) at 140 °C and 80 bar (1200 psi) total gas pressure. The
in situ intermediate was again an ester (ethyl formate), formed from the addition of ethanol
to the reaction mixture.273 This catalytic system also proved competent for direct CO2
reduction in the absence of ethanol, which proceeded via a homogeneous ruthenium-bound
formate species where formate coordination proved key to avoiding the thermodynamic
barrier for generation of a true formic acid intermediate. This adjustment yielded a
maximum TON of 442.274 While this was an encouraging improvement with respect to
TON and reaction simplicity, the system operated under acidic conditions that would be
incompatible with the bases often used in CO2 capture.
76
Rather than using alcohols additives to generate intermediate esters, CO2
hydrogenation can also proceed through formamide intermediates when it is performed in
the presence of amines. In this reaction, the CO2 reacts with the amine to form a carbamate,
after which metal-catalyzed dehydration gives the formamide. Final metal-catalyzed
hydrogenation yields methanol and regenerates the amine. Sanford et al. were the first to
incorporate this strategy, using a ruthenium-pincer catalyst (Figure 1.3.1, (3)) with
dimethylamine and potassium phosphate additives to generate methanol from CO2 (2.5 bar)
and H2 (50 bar) with TONs of up to 550.275 However, in order to maximize catalytic
activity, the authors had to devise a temperature-ramping method (95 °C to 155 °C) in order
to by-pass the dimethylcarbamate (DMC) pathway in the initial step and instead activate
the formation of their dimethylformamide (DMF) intermediate from dimethylformic acid
(DMFA). This resulted in a product mixture containing methanol, DMF, and DMFA, of
which the latter two were the major contributors (together accounting for more than twice
the amount of methanol present).275 The comparatively high yields combined with needing
only very low pressures of CO2 was incredibly promising, however, selectivity
improvements would be required before more widespread applications would be possible.
Further developments in the incorporation of amines for facilitating CO2
hydrogenation to methanol continued to center on ruthenium catalysts. Prakash et al.
applied polyamines to the exact reaction parameters established by Sanford and were able
to attain a top TON of 1060, which is the highest performing catalytic system to date.276
They have also advanced techniques for CO2 capture from air and amine recycling using
polyamines and biphasic solvents.277 Milstein and co-workers established a unique
approach that operates under very low CO2 pressures (< 3 bar) through the use of
77
Figure 1.3.1: Previous (1-5) and current metal catalysts for CO2 hydrogenation to
methanol via CO2 derivatives. OTf = triflate, OAc = acetate.
78
aminoethanes to capture the CO2 as a heterocycle in situ, followed by ruthenium-catalyzed
hydrogenation (albeit with TONs < 25).278 The hydrogenation pressures, however, are still
around 60 bar H2.
The literature concerning metal-catalyzed hydrogenation of CO2 to methanol is
dominated by precious-metal ruthenium catalysts, whereas the incorporation of base metals
is far more limited. In fact, there are only two examples of base-metal catalysts for this
transformation, which include one manganese-pincer complex279 and one cobalt-triphos
complex (Figure 1.3.1., (4) and (5)).280 While their structures are highly reminiscent of
their ruthenium counterparts, their performances are not, yielding maximum TONs of 36
and 78 for the production of methanol using manganese and cobalt, respectively (the
former proceeds via the amide pathway while the latter uses the ester option). There is
therefore room for significant improvement for precious-metal and base-metal catalysts
alike. For comparison, top-performing catalysts for CO2 hydrogenation to formic acid or
formate include Nozaki’s precious-metal iridium-pincer catalyst281 with TONs of over 3.5
million and a base-metal iron-pincer catalyst with TONs of greater than 46,000 (when
paired with a Lewis acid co-catalyst).282
Previous work in our laboratories and others has established the reactivity of iron-
pincer complexes of the type (RPNP)Fe(H)(CO) and (RPNMeP)Fe(H)(CO)(BH4) (RPNP =
N[CH2CH2(PR2)]2, RPNMeP = (CH3)N[CH2CH2(PR2)]2; R = iPr, Cy) in a multitude of
hydrogenative and dehydrogenative transformations, including CO2 hydrogenation to
formate (the example which was mentioned above),282 formic acid dehydrogenation,283 and
methanol dehydrogenation.284 In all cases the catalytic capabilities of these complexes far
exceed the performance of any other base-metal catalyst for their respective processes,
79
occasionally also surpassing the best precious-metal catalysts. It is therefore reasonable to
apply this class of compounds toward the problem of CO2 hydrogenation to methanol. In
addition, these catalysts have been successfully utilized by our laboratory in the N-
formylation of amines285 and the hydrogenation of formamides to amines and alcohols,286
so their catalytic aptitude with respect to the two separate steps of CO2 hydrogenation to
methanol in the presence of amines have also been established. This section summarizes
the ongoing investigation into the catalytic performance of the five-coordinate iron (II)
complex (iPrPNP)Fe(H)(CO) (1; Figure 1.3.2) with respect to the one-pot, two-step
synthesis of methanol from CO2 in the presence of morpholine (pathway (2) in Scheme
1.3.5). Even without being fully optimized, TONs of up to 360 have been attained for
methanol production from this reaction, representing an improvement of nearly half an
order of magnitude over other reported base-metal catalysts for CO2 hydrogenation.
1.3.2: Results and Discussion:
Initial studies in our laboratory with respect to 1-catalyzed CO2 hydrogenation to
methanol were conducted using dimethylamine287 due to its previous successful application
in this type of reaction.275 The production of methanol was not observed in the initial trials,
Figure 1.3.2: Structures of the iron-pincer complexes 1, 2, and 1-formate.
80
which contained 21 mmol dimethylamine in 6 mL dioxane with 5 mol% 1, 100 psi CO2,
and 900 psi H2 in a Parr reactor at 120 °C for 16 hours. However, a large amount of DMF
was generated and the competency of 1 with respect to DMF hydrogenation had already
been established, suggesting that the reaction should be able to proceed.287 Upon further
investigation, two major obstacles for obtaining methanol were identified. The first was
that the water released during the conversion of CO2 to the intermediate formamide was
significant enough to accelerate catalyst decomposition. The second was the inhibition of
the formamide hydrogenation step due to formation of (iPrPNHP)Fe(H)(CO)(COOH), or 1-
formate (Figure 1.3.2; confirmed using NMR experiments). In the presence of CO2 and
H2, complex 1 undergoes a 1,2-addition of H2 across the Fe-N bond, followed quickly by
CO2 insertion into the Fe-H bond to give 1-formate,282 which is a poor catalyst for amide
hydrogenation.287 The formate can be released from the iron complex using a combination
of a Lewis acid (lithium trifluoromethanesulfonate, or LiOTf) and a base (DBU), but if a
large concentration of CO2 is still present then 1 will also catalyze the competing reaction
of CO2 hydrogenation to formate (for which it can achieve several thousand TONs).282 In
order to solve these problems, a one-pot, two-step reaction sequence was adopted (Scheme
1.3.6). The first step (step A) involved N-formylation of the amine in the presence of 3 Å
molecular sieves to absorb the water generated in situ (the efficacy of which was confirmed
by NMR experiments). After this reaction was completed, the Parr reactor was partially
vented, brought into the glovebox, and opened. Then, the reaction solution was filtered to
remove the sieve dust as well as any solid alkylammonium carbamate. In the absence of a
CO2 headspace, this carbamate salt decomposes back to CO2 and amine, which would act
as an undesirable source of CO2 in the second step. After filtration, a known amount of
81
formamide (based on analysis of the step 1 solution) was loaded back into the Parr reactor
with more catalyst, along with LiOTf and DBU to regenerate 1 from any 1-formate formed
due to CO2 still dissolved in solution. This was followed by hydrogenation of the
formamide to give methanol and the free amine in the second step (step B).287
Despite these efforts, methanol was unable to be obtained with the use of
dimethylamine in the CO2 hydrogenation reaction.287 However, switching from
dimethylamine to morpholine immediately gave methanol TONs of greater than 250 using
the two-step procedure detailed above.287 Morpholine was recommended by its established
success as a substrate in both the N-formylation of amines and the hydrogenation of
amides.285,286 It was an attractive alternative to dimethylamine from a purification and
handling perspective due to the fact that it is a liquid at room temperature rather than a gas,
and its ring-constrained R groups had the potential to provide a steric advantage in
Scheme 1.3.6: Flow diagram for the one-pot, two-step reaction setup.
82
accessing hindered reaction intermediates. The morpholine-derived ammonium carbamate
salt was also a solid under ambient conditions rather than a liquid (like that of
dimethylamine), affording greater separation from the formamide before step B. Efforts to
re-integrate the two reaction steps described above and to optimize the reaction conditions
using morpholine have been the current focus of this work.
Initial experiments sought to determine whether it was possible to re-establish a
one-step hydrogenation reaction using a more cooperative amine (Tables A1.15 – A1.18).
The absence of molecular sieves in step A improved the amount of available 4-
formylmorpholine after step 1 (Table A1.15, entry 2), which was unsurprising as some
absorption of the liquid formamide by the sieves would be expected. The methanol yield,
however, was significantly reduced, from approximately 65% to 11% (Table A1.15, entry
3). Despite observed consumption of the 4-formylmorpholine starting material, no
methanol was detected when sieves were present during step B, which was attributed to
their complete absorption of the product (Table A1.15, entry 4). This indicated that
molecular sieves must be present for step A but removed before step B. Considering the
possibility of using a different drying agent, LiOTf and DBU were also screened for their
compatibility between steps. Results indicated that while LiOTf was integral to the
performance of the second step, it detrimentally affected the yield of the first (Table
A1.16). The effect of DBU was found to be concentration-dependent (Table A1.17). The
presence of DBU alone did not harm either step (entry 3), however, if the overall amount
of DBU in solution was doubled, the yield of methanol began to decrease (entry 4). DBU
also caused practical issues with GC-FID analysis of the step A solution due to its affinity
for the column stationary phase, causing its signal to bleed into other parts of the spectrum
83
and giving the appearance of reaction yield enhancement by its presence. Finally,
morphylammonium morphylcarbamate was independently synthesized (confirmed by
crystal structure, see Figure A1.52), isolated, and added to step B to see if the reaction
could proceed favorably in its presence (Table A1.18). It was found to have a significant
detrimental effect on methanol production. The combination of these results confirmed the
necessity of maintaining the two-step procedure, despite the amenability of morpholine
toward methanol production.
Two-step reaction optimization studies began with screening solvents (Table
A1.19), which had very little effect on methanol production, except for in the case of ethyl
acetate where an unidentified side product was observed by GC-FID. However, it was
indicated that tetrahydrofuran (THF) would be an acceptable substitute for 1,4-dioxane
(Table 1.3.1, entry 1 versus entry 2), which is an easier solvent to handle in terms of drying
procedures. It was therefore the solvent of choice in all future experiments. The methanol
yields from the solvent screen also suggested that a higher loading of 4-formylmorpholine
would be required in step B in order to truly asses the catalytic limitations of 1, since step
B essentially went to completion in all cases except ethyl acetate. Due to the fact that some
of the product is lost to the molecular sieves as well as during the filtration after step A, a
means of increasing that yield had to be devised. Increasing the ratio of CO2 in the gas
mixture (Table A1.20) afforded a moderate gain in product yield. Rather than continue to
increase CO2 levels, which is disadvantageous for step B, the overall reaction pressure was
increased from 1000 psi to 1400 psi and the formamide production was more than doubled
as a result (Table 1.3.1, entry 2A versus entry 3A). The formamide loading in step B could
then be increased from 1.25 mmol to 4 mmol, providing a larger window with which to
84
fine-tune improvements. The amounts of the LiOTf and DBU additives were varied to
confirm that their concentrations were sufficient to handle the effects of a higher substrate
loading in step B, and it was found that no changes were needed (Table A1.21). It was also
confirmed that loading fresh catalyst solution in step B was necessary to acquire any
detectable methanol.
Entryd Substrate
(mmol)
CO2/H2
(psi)
Catalyst Solvent Temp.
(°C)
FM
(mmol)
CH3OH
(mmol)
TONc
1A Morph.
(21)
100/900 1 1,4-
dioxane
120 4.30 - 286
1B FM
(1.25)
0/900 1 1,4-
dioxane
120 0.03 1.14 114
2A Morph.
(21)
100/900 1 THF 120 4.42 - 294
2B FM
(1.25)
0/900 1 THF 120 0.01 1.13 113
3A Morph.
(21)
200/1200 1 THF 120 9.24 - 616
3B FM
(4.0)b
0/1200 1 THF 120 0.05 3.40 340
4A Morph.
(21)
200/1200 2 THF 120 3.20 - 213
4B FM
(4.0)b
0/1200 - - - - - -
5A Morph.
(21)
200/1200 1 THF 100 16.76 - 1117
5B FM
(4.0)b
0/1200 1 THF 100 0.01 3.60 360
Table 1.3.1: Select Reaction Optimization Data for CO2 Hydrogenation to Methanol in
the Presence of Morpholine.a
aReaction conditions: Step A: 21 mmol morpholine (morph.), 2.02 g molecular sieves, and 15 µmol
catalyst 1 or 2 in 8 mL solvent with X psi CO2 and Y psi H2 at Z °C for 16 hours. Step B: 1.25 mmol
formyl morpholine (FM) from step A, 1 mmol LiOTf, 2.5 mmol DBU in a total volume of 5 mL with Y
psi H2 at Z °C for 16 hours. b4.0 mmol FM in step B. cTON = turnover number for FM in step A and
CH3OH in step B as determined by GC-FID analysis. dAverage of two trials.
85
Due to the necessity of a two-step reaction process, there was an opportunity for
employing different catalysts in steps A and B in order to maximize the productivity at
each stage. As mentioned briefly in the introduction, previous work in our laboratory has
used the iron pincer complexes 1 and 2 (Figure 1.3.2) for the N-formylation of amines,
where the performance of 2 exceeded that of 1 (attributed in part to its greater stability in
the presence of water). As water sensitivity has been identified as a limiting factor for
complex 1, complex 2 was utilized for N-formylation in first step, with the intention of
applying complex 1 to amide hydrogenation in the second step. However, it was found that
under these reaction conditions, the formamide yield was only a third of what it was when
complex 1 was used as the catalyst (Table 1.3.1, entry 3 versus entry 4). Further
investigations indicated that the presence of the molecular sieves interferes with the
catalytic ability of 2 (Table A1.22), which was tentatively ascribed to their slight acidity.
It is possible that this is also a contributing factor in the behavior of complex 1, as there is
a similar significant increase in formamide yield for step A in the absence of sieves (Table
A1.15) that was originally explained as a matter of product absorption by the additive.
More experiments will be required before a definitive understanding of this phenomenon
can be reached.
The final set of optimizations with respect to temperature found that a 20 °C
decrease nearly doubled the production of 4-formylmorpholine in step A and had no
detrimental effect on step B (Table 1.3.1, entry 3 versus entry 5). Lowering the temperature
further (80 °C) resulted in a slight decrease in yield (Table A1.23). This trend is attributed
to a balance between having a high enough temperature to favor the reaction kinetically
while avoiding thermally-induced catalyst degradation. The TON of 360 achieved for
86
methanol production in step B under this milder set of temperature conditions represents
the highest activity of any base-metal catalyst for CO2 hydrogenation to methanol by nearly
half an order of magnitude.
1.3.4: Summary
In this section, the iron-pincer complex (iPrPNP)Fe(H)(CO) (complex 1) was
introduced, which belongs to a class of iron catalysts that display remarkable activity with
respect to hydrogenation and dehydrogenation reactions, especially those incorporating
CO2 (a particularly inactive substrate). The initial studies presented here revealed its
proficiency for the hydrogenation of CO2 to methanol in the presence of amines, and it far
outperformed any other base-metal catalyst for this transformation. Its best TON of 360
was also highly competitive with the reported activities of its precious-metal counterparts.
The promising catalytic capability of 1 therefore represents significant advancement
toward the goal of the more widespread utilization of CO2 as a carbon feedstock, since by
using CO2 as a direct source for methanol it also indirectly sources all of the materials that
can be made from methanol. The incorporation of abundant, less expensive base metals in
the catalysts for these processes also assists in maintaining the economic benefit and long-
term sustainability of transitioning to such a C1 carbon source. However, further
experiments will need to be conducted before the upper catalytic limit of complex 1 can be
determined. Unfortunately, the water sensitivity of 1 is a restrictive factor, and devising a
means of bypassing this without sacrificing catalytic performance is of the utmost
importance. The use of molecular sieves, while effective, prohibits the application of
structurally similar but more highly active catalytic systems for the first step of the
hydrogenation process, which implies a previously unrecognized influence on chemical
87
reactivity that needs to be further elucidated. The investigation of other drying agents or
making adjustments to the step A processing methods (i.e. add sieves after step A is
complete and filter just before step B) may provide alternative (albeit more costly or
energy-intensive) solutions. Another factor to consider is the necessity of using one of the
highest total pressures reported for this hydrogenation reaction in order to maximize
conversion. Following the ester rather than the amide pathway through the use of alcohol
additives in place of amines may be more kinetically accessible and help to alleviate the
need for such forcing conditions. This would also avoid the formation of an ammonium
carbamate, which can serve as a masked CO2 source that is detrimental to step B catalysis.
The impressive performance of 1 in the face of these challenges, however, suggests at a
much greater potential once they have been overcome.
88
1.4: Chapter Summary
In this chapter, the catalytic competency of nickel and iron complexes with respect
to the activation of CO2 and CO2 analogues and their subsequent conversion into valuable
commercial chemicals was explored. These transformations were enacted using either
alkenes or dihydrogen as highly reactive co-reagents in order to make the reactions more
energetically favorable. In the first two sections, it was found that the encouraging
stoichiometric coupling performances of nickel promotors in CO2/alkene/CO and
isocyanate/alkene reactions did not translate to product selectivity, which would need to be
overcome (likely via systematic ligand variation) before they could be extended to catalytic
investigations. The last section of this chapter detailed ongoing studies regarding the
promising ability of the iron pincer complex (iPrPNP)Fe(H)(CO) (complex 1) in the
hydrogenation of CO2 to methanol. Even without final optimization of the reaction
conditions, complex 1 achieves TONs that are nearly five times those of any other base-
metal catalyst and that match or surpass those of most of its precious metal competition for
the production of methanol. This performance therefore denotes a significant improvement
in catalytic capability, the importance of which rests in the opportunity the realization of
this transformation represents for sourcing not only methanol but the multitude of products
it engenders from a cheap, abundant, and renewable carbon source. That being said,
addressing the issues of the water-sensitivity of 1 (necessitating a two-step reaction
sequence) and the use of reaction pressures on the higher end of those reported to achieve
the best TONs will likely be needed before its more widespread application is realized.
Overall, while the results presented in this chapter are modest, the reactions they
address are inherently challenging. For example, nickel-catalyzed production of acrylate
89
from CO2/ethylene coupling has only recently been attained after over 40 years of study,
with TONs on the order of ~100.117 Understanding the role of the base, Lewis acid, and
zinc additives that made this catalysis possible will likely impact the field as a whole and
help to advance analogous nickel-based reactions, like those discussed herein. With only a
handful of examples that have achieved at best 1000 turnovers, the metal-catalyzed
hydrogenation of CO2 to methanol is a similarly non-trivial transformation. A pivotal issue
is the beneficial impact of high CO2 content on the first half of the reaction but its inhibition
of the final hydrogenation step. Evolving methods of using low CO2 pressures but still
maintaining high overall yields will likely be key to making this reaction more competitive
with other CO2 hydrogenation/dehydrogenation cycles.
90
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A Practical Approach from CO2 and Epoxides to Methanol and Diols,” Angew. Chem. Int. Ed., 2012, 51,
13041-13045.
273 Wesselbaum, S.; vom Stein, T.; Klankermayer, J.; Leitner, W. “Hydrogenation of Carbon Dioxide to
Methanol by Using a Homogeneous Ruthenium-Phosphine Catalyst,” Angew. Chem. Int. Ed., 2012, 51, 7499-
7502.
274 Wesselbaum, S.; Moha, V.; Meuresch, M.; Brosinski, S.; Thenert, K.M.; Kothe, J.; vom Stein, T.; Englert,
U.; Hoelscher, M.; Klankermayer, J.; Leitner, W. “Hydrogenation of carbon dioxide to methanol using a
homogeneous ruthenium-Triphos catalyst: from mechanistic investigations to multiphase catalysis,” Chem.
Sci., 2015, 6, 693-704.
275 Rezayee, N.M.; Huff, C.A.; Sanford, M.S. “Tandem Amine and Ruthenium-Catalyzed Hydrogenation of
CO2 to Methanol,” J. Am. Chem. Soc., 2015, 137, 1028-1031.
276 Kothandaraman, J.; Goeppert, A.; Czaun, M.; Olah, G.A.; Prakash, G.K.S. “Conversion of CO2 from Air
into Methanol Using a Polyamine and a Homogeneous Ruthenium Catalyst,” J. Am. Chem. Soc., 2016, 138,
778-781.
277 Kar, S.; Sen, R.; Goeppert, A.; Prakash, G.K.S. “Integrative CO2 Capture and Hydrogenation to Methanol
with Reusable Catalyst and Amine: Toward a Carbon-Neutral Methanol Economy,” J. Am. Chem. Soc., 2018,
140, 1580-1583.
278 Khusnutdinova, J.R.; Garg, J.A.; Milstein, D. “Combining Low-Pressure CO2 Capture and Hydrogenation
To Form Methanol,” ACS Catal., 2015, 5, 2416-2422.
279 Kar, S.; Goeppert, A.; Kothandaraman, J.; Prakash, G.K.S. “Manganese-Catalyzed Sequential
Hydrogenation of CO2 to Methanol via Formamide,” ACS Catal., 2017, 7, 6347-6351.
280 Schneidewind, J.; Adam, R.; Baumann, W.; Jackstell, R.; Beller, M. “Low-Temperature Hydrogenation
of Carbon Dioxide to Methanol with a Homogeneous Cobalt Catalyst,” Angew. Chem. Int. Ed., 2017, 56,
1890-1893.
281 Tanaka, R.; Yamashita, M.; Nozaki, K. “Catalytic Hydrogenation of Carbon Dioxide Using Ir(III)-Pincer
Complexes,” J. Am. Chem. Soc., 2009, 131, 14168-14169.
282 Zhang, Y.; MacIntosh, A.D.; Wong, J.L.; Bielinski, E.A.; Williard, P.G.; Mercado, B.Q.; Hazari, N.;
Bernskoetter, W.H. “Iron catalyzed CO2 hydrogenation to formate enhanced by Lewis acid co-catalysts,”
Chem. Sci., 2015, 6, 4291-4299.
283 Bielinski, E.A.; Lagaditis, P.O.; Zhang, Y.; Mercado, B.Q.; Würtele, C.; Bernskoetter, W.H.; Hazari, N.;
Schneider, S. “Lewis Acid-Assisted Formic Acid Dehydrogenation Using a Pincer-Supported Iron Catalyst,”
J. Am. Chem. Soc., 2014, 136, 10234-10237.
284 Bielinski, E.A.; Förster, M.; Zhang, Y.; Bernskoetter, W.H.; Hazari, N.; Holthausen, M.C. “Base-Free
Methanol Dehydrogenation Using a Pincer-Supported Iron Compound and Lewis Acid Co-catalyst,” ACS
Catal., 2015, 2, 2404-2415.
107
285 Jayarathne, U.; Hazari, N.; Bernskoetter, W.H. “Selective Iron-Catalyzed N-Formylation of Amines using
Dihydrogen and Carbon Dioxide,” ACS Catal., 2018, 8, 1338-1345.
286 Jayarathne, U.; Zhang, Y.; Hazari, N.; Bernskoetter, W.H. “Selective Iron-Catalyzed Deaminative
Hydrogenation of Amides,” Organometallics, 2017, 36, 409-416.
287 Unpublished work by Dr. Yuanyuan Zhang
108
Chapter 2:
The Application of an Iron Catalyst in the Transformation of Methanol into Formamides
and Ureas via Acceportless Dehydrogenation Reactions
Reprinted (adapted) with permission from Lane, E.M.; Uttley, K.B.; Hazari, N.;
Bernskoetter, W.H. “Iron-Catalyzed Amide Formation from the Dehydrogenative
Coupling of Alcohols and Secondary Amines,” Organometallics, 2017, 36, 2020-2025.
Copyright 2017 American Chemical Society.
Reprinted (adapted) with permission from Lane, E.M.; Hazari, N.; Bernskoetter, W.H.
“Iron-Catalyzed Urea Synthesis: Dehydrogenative Coupling of Methanol and Amines,”
Chem. Sci., 2018, 9, 4003-4008. Published by the Royal Society of Chemistry.
109
2.1: Introduction: Methanol as a Carbon Feedstock
As mentioned in Chapter 1, methanol is currently produced on an industrial scale
from a synthesis gas mixture (syngas; H2/CO2/CO) using alumina-supported mixed copper
and zinc oxide heterogeneous catalysts at elevated temperature and pressure (Scheme 2.1.1,
reaction (1)).1,2 While most syngas is derived from coal and natural gas,3 it can be
renewably sourced from biomass (via gasification),4-6 albeit at lower yields than its
nonrenewable counterparts.7,8 As of 2016, global methanol demand had reached 80 million
metric tons and since then has continued to rise.9,10 Although methanol is often used as a
solvent, >95% of methanol produced is subsequently converted into other valuable
chemicals,9 making it an abundant and potentially renewable carbon feedstock (Scheme
2.1.2).3 As raw material feedstocks typically constitute 60-70% of manufacturing
expenses,3 the low cost of methanol provides a significant economic advantage in this
respect. Formaldehyde accounts for the most substantial percentage of methanol demand
(27%),9 however, its value is largely based on its role as a highly active intermediate on
the pathway from methanol to more functionalized species. For example, further
transformation of formaldehyde yields high-demand products like resins for building
Scheme 2.1.1: (1) Industrial methanol production from syngas and (2) industrial
formaldehyde production from methanol.
110
materials, industrial and automotive plastics, insulation, and pharmaceuticals.3,11 Industrial
production of formaldehyde is achieved through methanol oxidation/dehydrogenation
using either a silver or iron-molybdenum oxide heterogeneous catalyst in a fixed-bed
reactor at high temperatures and pressures (Scheme 2.1.1, reaction (2)).12-14 These
conditions are appropriate for formaldehyde, but they are incompatible with the additional
functionalization necessary to give formaldehyde-derived desirable products. This is a
common issue amongst many of the first generation methanol-sourced compounds in
Scheme 2.1.2, and as a result, commodity chemical production from methanol is relegated
to a multi-step process. An alternative method to achieve the goal of methanol
functionalization is to instead use metal-catalyzed dehydrogenative coupling under milder
conditions to convert methanol to the end product (via a formaldehyde intermediate) in one
pot, turning it into a more effective C1 carbon source.
Scheme 2.1.2: Some of the major chemical species derived from methanol and their
uses.
111
Metal-catalyzed alcohol dehydrogenation effectively transforms an alcohol into a
more reactive aldehyde or ketone species without stoichiometric amounts of harsh, toxic,
or waste-producing oxidizing agents.15,16 The absence of such reagents opens the door to a
host of previously unavailable manipulations, as homogeneous organometallic catalysts
often suffer ligand degradation under oxidative conditions and therefore their applicability
to such processes was previously limited.17 The first known example of metal-promoted
alcohol dehydrogenation was introduced in the late 1800’s by Guerbet and was enacted
under harshly basic and high temperature conditions (130-180 °C) with heterogeneous
catalysts.18 Significant advancement in the field has occurred since then and is extensively
covered in several reviews.16-21 The concept of dehydrogenation as a form of alcohol
activation is a much more recent development, however, and while a host of examples now
exist, it was the early work by the groups of Fujita and Yamaguchi,22 Ramón and Yus,23
and Williams24,25 that generalized the method for a variety of transformations and formally
established the idea. By this process, a nucleophilic alcohol is transformed into an
electrophilic aldehyde or ketone, priming it for subsequent reaction with a different
nucleophile and further functionalization.18,23 In this manner a variety of bond-forming
reactions (such as C-C, C-N, and C-O)16,18,24,24,82 can be enacted to yield valuable
functionalized products from mild, cheap, and easy to handle starting materials. Alcohol
dehydrogenation followed by tandem coupling therefore accesses a diverse assortment of
useful products in a more environmental and economical manner.
In traditional alcohol dehydrogenation, the product of the reaction of the generated
aldehyde/ketone with other substrates also serves as an internal hydrogen acceptor to re-
activate the hydrogenated catalyst.23 However, this interferes with the generation of
112
unsaturated products. One solution to this problem is to use a sacrificial hydrogen
acceptor26 like an alkene (e.g. tert-butyl ethylene)27 or a ketone (e.g. acetone),28 but these
additives result in additional reaction cost and waste generation.17 It is therefore
advantageous to incorporate a catalyst that can release the hydrogen it removes from the
alcohol as dihydrogen, referred to as acceptorless alcohol dehydrogenation. This is
typically accomplished using higher reaction temperatures, which enhance the kinetics and
thermodynamics of the reaction by speeding it up and by providing an entropic driving
force via the production of two species rather than one (respectively). In addition,
performing the dehydrogenation reaction at reflux temperatures effectively removes the
dihydrogen from solution, further favoring the process.17,29 However, even this byproduct
has value, as reversible dihydrogen generation from methanol has potential alternative
energy applications and there have been extensive investigations into the use of methanol
as a hydrogen storage material due to its high energy density (see chapter 1 for more
information).
Methanol dehydrogenation presents an additional challenge due to a greater
energetic demand for this process (ΔH = +84 kJ/mol) compared to that for higher alcohols
(e.g. ethanol ΔH = +68 kJ/mol)14,30,31 The formaldehyde intermediate initially produced is
also very unstable and will readily decompose to CO and H2 if it is not trapped effectively.65
As a result, the number of catalysts capable of methanol activation and functionalization is
limited, and those incorporating cheap and abundant base metals are even rarer.62,65,82 In
this chapter, a five-coordinate iron (II) PNP pincer complex for the catalytic acceptorless
dehydrogenative coupling of methanol with amines to give commercially valuable
formamides and ureas is introduced. Its impressive activity in this respect matches and
113
in the case of amidation even exceeds the performance of precious-metal catalysts,
representing significant advancement in this approach to methanol transformation.
114
2.2: Iron-Catalyzed Amide Formation from the Dehydrogenative Coupling of
Alcohols and Secondary Amines
2.2.1: Introduction
Amides are key functional groups in organic synthesis, biochemistry, and
pharmaceuticals.32 For example, it is estimated that approximately 25% of drug
molecules contain amides.33 Current methods for large-scale amide synthesis
typically involve the reaction of carboxylic acids33,34 or more commonly their
activated derivatives (such as acid chlorides, acid anhydrides, and acyl azides)35-37
with stoichiometric amounts of amines (Scheme 2.2.1), and additional inorganic or
organic promoters (e.g. bases, boronic acids, or borate esters) are often used. As a
result, these procedures are not atom efficient, they generate considerable chemical
waste, and they use relatively expensive starting materials.32 The dehydrogenative
coupling of alcohols and amines to form amides (with dihydrogen as the only
byproduct, Scheme 2.2.2) offers an atom-economical and more environmentally
Scheme 2.2.1: Metal-free methods for amide production.
115
benign alternative to obtaining these materials. Although coinage metals such as
silver38 and gold39-42 have demonstrated some effectiveness as heterogeneous
catalysts for dehydrogenative amide synthesis, these systems normally require high
catalyst loadings, harsh reaction conditions, long reaction times, and have limited
substrate scopes. The sole heterogenous copper example exhibits a slightly wider
scope but otherwise suffers from similar issues, as well as the necessity of an
oxidatively harsh peroxide additive.43 Homogenous transition metal catalysts
incorporating rhenium,44 rhodium,45 and (most extensively) ruthenium have shown
greater activity.46-63 The most prominent example, a ruthenium-pincer catalyst
published in Science in 2007 by Milstein and co-workers, achieved TONs of nearly
1000 for the dehydrogenative coupling of primary amines with sterically accessible
alcohols (Figure 2.2.1).47 This spurred the development of a host of other
predominantly precious-metal catalysts for dehydrogenative amide synthesis, but
nearly all suffered from the limitations of requiring an added base, needing a
stoichiometric amount of a hydrogen acceptor, giving low TONs (<100), or a
combination of these factors. Notably, most existing catalysts perform best with
Scheme 2.2.2: Metal-promoted pathways for dehydrogenative amidation.
116
primary amine substrates, and there are only a limited number of catalysts (all but
two are ruthenium-based) that are compatible with a range of secondary amine
substrates.51,52,54,56,58-60,62,64,65 To date, only two of these catalysts have achieved
TONs greater than 50 for the formation of tertiary amides under base-free and/or
hydrogen acceptor-free conditions (Figure 2.2.2).56,59 Furthermore, the development
of promotors incorporating cheaper, less toxic metals is lacking, as only three base-
metal catalysts (one copper and two manganese) have been described for
dehydrogenative amidation (Figure 2.2.1).64-66 While these recent studies
demonstrate some promise of base-metal-mediated coupling of amines and alcohols,
Figure 2.2.1: A top-performing ruthenium catalyst (a) and all previous base-metal
catalysts (b) for dehydrogenative amide formation from amines and alcohols.
117
they exhibit limited tertiary amide production (TONs <50) and leave open the
possibility for substantial improvement.
As previously mentioned, smaller alcohols like methanol have to overcome a
larger activation energy for dehydrogenation, which makes them incompatible with
many catalysts.31 However, formamide production from the dehydrogenative
coupling of methanol and amines is a crucial target. Formamides not only appear in
agrochemicals, dyes, and fragrances,67,68 but they serve as reagents69 (e.g. in the
Vilsmeier-Haack reaction)70 and are key organic intermediates, such as in the
production of formamidines71 and isocyanides.72,73 They are also important amine
protecting groups in protein biosynthesis.74-76 Perhaps their most valuable trait,
however, is their high biological activity, which makes them integral to various
pharmaceuticals.62,77 Formamides are typically synthesized from amines and excess
formylating agents (chloral, phenyl formate, formic acid), which are
environmentally hazardous, generate copious waste, and display a small range of
utility.78,79 Most of the few existing methods for using methanol as a “green” C1
building block in place of these hazardous formylating reagents in formamide
syntheses require harsh conditions, long reaction times, and show limited scope.41-
43,62 Of the many previously reported homogeneous catalysts for amide formation
from alcohols and amines, only four (three ruthenium-based and one manganese-
based) were capable of formamide production (Scheme 2.2.2, R=H). All four
exhibited low activities, three58,62,65 with TONs ≤ 50 (Figure 2.2.2) and one
ruthenium catalyst80 with TONs < 100 but which required a base additive, displayed
118
poor selectivity, and was only investigated with one amine substrate. Therefore, a
catalyst capable of the dehydrogenative coupling of methanol and amines in the
Figure 2.2.2: Efficient ruthenium, manganese, and iron catalysts for the dehydrogenative
coupling of alcohols and 2° amines.
119
absence of other reagents is both desirable and rare, especially one incorporating a
base metal.
Several independent research endeavours (including within our laboratories) have
reported the successful application of iron-pincer catalysts of the type
(RPNP)Fe(H)(CO) (RPNP = N[CH2CH2(PR2)]2, R = iPr, Cy) in the dehydrogenation
of alcohols to esters,81 ketones,81 or carbon dioxide.82,83 Recently, we also described
the use of the five-coordinate iron (II) complex (iPrPNP)Fe(H)(CO) (1) in the
catalytic base-free hydrogenation of amides,84 which suggests it may have the
potential to catalyze the microscopic reverse reaction (the dehydrogenative
formation of amides). The following discussion summarizes the impressive ability
of 1 to catalyze the dehydrogenation of alcohols in the presence of secondary amines
to selectively produce amides in the absence of secondary hydrogen acceptors or
base promotors (Figure 2.2.2). Complex 1 is the most productive base-metal catalyst
yet developed for intermolecular dehydrogenative amide synthesis and gives TONs
that exceed those previously described for base or precious metals with respect to
tertiary amide formation.85
2.2.2: Results and Discussion
Catalytic Studies
Due to the success of 4-formylmorpholine as a substrate in prior amide
hydrogenation experiments,84 investigations began with attempts to couple
morpholine and methanol in the presence of 1. An initial experiment was performed
using a 4:1 amine:alcohol ratio and 0.1 mol% of 1 in 10 mL of dioxane at 80 °C for
8 hours. Despite the lack of optimization, this reaction yielded a promising TON of
120
297 for the formation of 4-formylmorpholine. This TON is already several times
that exhibited by any other base-metal catalyst for this reaction. Using this reaction
as a benchmark, the solvent, reaction volume, time, temperature, and substrate ratio
were all varied to generate an optimized set of reaction conditions (Tables A2.1-
A2.5). Using these optimized conditions (illustrated in Table 2.2.1), a TON of 503
was achieved for the formation of 4-formylmorpholine (entry 3), which rivals the
highest TONs reported for ruthenium catalysts in dehydrogenative amidation using
secondary amines and far exceeds those of base-metal catalysts. The TON was
confirmed by NMR analysis of the 4-formylmorpholine product and by collection
of the dihydrogen produced (86% of expected H2 was collected). Encouraged by the
activity of 1 with traditionally more difficult secondary amines, the substrate scope
of dehydrogenative amidation was further explored. Overall, 1 is extremely effective
at amide production using cyclic secondary amines (entries 1-3), with a maximum
TON of 600 observed using 1,2,3,4-tetrahydroisoquinoline. However, the catalytic
activity appears strongly correlated to the size of the substituents on the amine. For
example, in comparing sterically restricted cyclic amines to substrates with even
small linear substituents, such as ethyl, there exists a significant drop in TON down
to approximately 200 (entry 4). This effect is further demonstrated as the amine
substituents increase in size (entries 5-7), with only trace conversion observed for
the phenyl and isopropyl derivatives. Attempts to generate diamides starting from
cyclic amides (entries 8 and 9) were also unsuccessful, although previous studies
from our laboratories suggest that this may be due to a competing 1,2-addition
reaction between secondary amides and 1.84
121
Entry Amine TON
1
600b
2
564b
3
503b
4 (CH3CH2)2NH 213
5 (C6H5CH2)2NH 126
6 (C6H5)2NH <10
7 [(CH3)2CH]2NH <10
8
0b
9
0b,c
Table 2.2.1: Dehydrogenative Amidation of
Methanol Catalyzed by 1.a
a Reaction conditions: 3 µmol catalyst (0.1 mol%),
3 mmol alcohol, and 12 mmol amine in 5 mL THF
at 80 °C for 8 hrs. Each entry is an average of two
trials and the TON was determined by GC analysis
of amide production unless otherwise indicated. bDetermined by 1H NMR spectroscopy.
cOnly one
trial.
122
Initial studies into extending the scope of amines to include primary amine
substrates have indicated a propensity for 1 to further substitute the initially formed
formamides to yield ureas, which is the subject of the next section in this chapter.
Despite the aforementioned steric limitations on the amine substrates, the ability of
1 to dehydrogenatively activate methanol and thereby catalyze the production of a
range of formamides with TONs of up to 600 in the absence of any additives is
unrivalled by base- and precious-metal catalysts alike. These results therefore
represent a significant step forward with respect to the use of methanol as a carbon
feedstock for more complex and commercially valuable organic species via
homogeneously catalyzed dehydrogenative coupling reactions.
The success of 1 in dehydrogenating a challenging substrate like methanol
motivated an investigation into the alcohol substrate scope for amidation (Table
2.2.2). Despite the fact that the primary focus of this discussion has been on the use
of methanol as a carbon source, access to a suite of alcohols capable of this
transformation would maintain the advantages of this synthetic approach while
further diversifying the product pool. Similar to the observations with amines, the
performance of 1 is highly sensitive to the steric profile of the alcohol (entries 1-3).
This effect on catalytic activity remains even when the larger substituent is slightly
removed from the dehydrogenation site (entries 4 and 5). The absence of catalysis
in the presence of 2,2,2-trifluoroethanol (entry 6) is likely due to its relative acidity
and its electron-withdrawing substituent. Interestingly, methanol far outperforms the
other alcohols explored. When this observation is combined with the results on the
effect of sterics on the amine substrate, it indicates that there is a strong steric effect
123
at both sides of the product carbonyl moiety. To some extent these results are
unsurprising, as large groups would potentially interfere with metal-catalyzed
dehydrogenation steps in which the substrates or intermediates must pass the
isopropyl substituents of the pincer ligand. Although the overall substrate scope
exhibited by 1 remains quite sterically limited, the high productivity observed for
small substrates suggests that the inherent activity of the iron center is significant
and could perhaps be leveraged by further attenuation of the ancillary ligand
environment.
Mechanistic Investigations
Entry Alcohol TON
1 CH3OH 503
2 CH3CH2OH 50
3 CH3(CH2)5OH 13
4 C6H5(CH2)2OH 10
5 C6H5CH2OH 10
6 CF3CH2OH 0b
Table 2.2.2: Alcohol Dehydrogenation in the
Presence of Morpholine Catalyzed by 1.a
aReaction conditions: 3 µmol catalyst (0.1 mol%),
3 mmol alcohol, and 12 mmol amine in 5 mL THF
at 80 °C for 8 hrs. Each entry is an average of two
trials and TON was determined by GC analysis of
amide production unless otherwise indicated. bDetermined by 1H NMR spectroscopy.
124
Recent investigations of transition-metal-catalyzed dehydrogenative amidation
have nearly all proposed the same general pathway for amide formation (Scheme
2.2.1).47,49,61 Each begins with the dehydrogenation of the alcohol to the
corresponding aldehyde followed by trapping either with amine or an additional
equivalent of alcohol to generate a hemiaminal or hemiacetal intermediate,
respectively. Dehydrogenation of an intermediate hemiaminal would directly afford
amide, while oxidation of the hemiacetal would yield an ester which could then
produce amide via a transamidation process. However, the exact mechanisms for
dehydrogenative amidation are likely catalyst and system dependant. For example,
many ruthenium catalysts employ base additives to achieve optimum performance,
which can potentially serve many different roles. These include enhancing
coordination of the alcohol to the metal,50 facilitating amine-to-aldehyde proton
transfer, obviating hemiaminal formation,61 or promoting ester transamidation.54
Prior studies within our laboratories demonstrate that 1 is quite capable of
promoting ester formation, achieving a TON of 107 over just 12 minutes for the
methanol to methylformate conversion under conditions similar to those used in this
work for amidation.82 Additional amidation experiments using catalyst 1 and
morpholine with methyl formate in place of alcohol yielded substantial quantities 4-
formylmorpholine (Scheme 2.2.3, Table A2.7).86 These observation suggest that the
lower pathway illustrated in Scheme 2.2.1 is at least viable for 1-catalyzed
amidation. However, further mechanistic experiments substituting benzaldehyde
for alcohol during the amidation process similarly afforded the corresponding amide
product (Scheme 2.2.3, Table A2.8), which also validates the hemiaminal-dependent
125
pathway for amide production.86 Since these preliminary investigations indicate that
both of the Scheme 2.2.1 pathways are competent for 1, it is clear that a more
exacting and extensive mechanistic undertaking will be require to elucidate the
precise path of 1-catalyzed amidation.
During the course of the mechanistic and catalytic experiments described above,
it was noted that when the catalytic experiments were assembled, the solution of 1
and amine immediately changed color from deep red to yellow-orange upon addition
of the alcohol substrate at room temperature. When the solution was heated during
catalysis, the color would quickly revert back to a shade of deep red resembling its
original hue. In order to ascertain any mechanistic implications that this color change
might entail, in situ NMR spectroscopic studies using morpholine and methanol
were performed. Multiple experiments showed that the color change and
corresponding changes to the 31P and 1H NMR spectra were exclusively dependent
on the addition of alcohol (Figure A2.1), since treatment of 1 with only morpholine
produced no observable reaction. The reaction of 1 with methanol was examined
using variable-temperature NMR spectroscopy, which indicated an equilibrium
between 1 and the iron-methoxy species (iPrPN(H)P)Fe(H)(OCH3)(CO) (2) formed
Scheme 2.2.3: Iron-catalyzed coupling of benzaldehyde or methyl formate with
morpholine.
126
by the addition of methanol across the Fe-N bond. At all observed temperatures
between 22 and 70 °C, only one 31P NMR resonance was seen, although its chemical
shift changes with temperature (Figure 2.2.3). This is indicative of a rapid
equilibration between 1 and 2 on the timescale of the NMR experiment, and the
position of the resonance represents an average value for the chemical shifts of the
two species present in solution. The change in the 31P NMR chemical shift suggests
that 1 is favored at high temperatures and 2 is preferred at lower temperatures, as
expected for a bimolecular to unimolecular reaction. The identity of 2 was confirmed
using 13CH3OH labeling and low-temperature NMR studies, which slowed the
Figure 2.2.3: Variable temperature 31P NMR
spectra indicate an equilibrium between 1 and 2
in the presence of 6 equiv. of methanol. Sample
dissolved in benzene-d6.
1
2 22 °C
50 °C
60 °C
70 °C
127
exchange rate between free and iron-activated methanol. At -80 °C, two distinct
peaks for the bound and free 13CH3 fragment were observed in the 13C {1H} spectrum
and a sharp resonance for 2 was seen at 93.2 ppm in the 31P NMR spectrum (Figure
A2.2). A proton-coupled 13C spectrum showed that the labeled carbon is still
attached to three protons, supporting the presence of an iron-methoxy species
(Figure A2.2). The exact position of the 31P{1H} NMR peak for 2 depends both on
temperature and the number of equivalents of methanol due to the equilibrium
described above.
The formation 2 via deprotonation of methanol is likely an off-cycle reaction that
competes with catalytic amidation. As mentioned above and illustrated in Scheme
2.2.1, the dehydrogenation of an alcohol to an intermediate aldehyde has been
implicated in most related amidation processes.47,49,61 Our prior experimental and
computational studies of 1-catalyzed methanol dehydrogenation have found that this
reaction proceeds via a concerted transfer of the hydrogenation atoms associated
with the methanol O-H and C-H bonds. Dehydrogenation pathways involving
intermediates analogous to complex 2 were found to be significantly higher in
energy.82 Thus, while formation of 2 may have some kinetic relevance to catalytic
amidation, it is likely not directly on its pathway and the high-temperature catalytic
conditions presumably minimize the significance of 2 in amide formation.
2.2.3: Summary
The five-coordinate iron complex 1 exhibits the highest productivity yet
reported for base-metal-catalyzed intermolecular dehydrogenative coupling of
alcohols and amines to amides. This iron-mediated amidation displays an intriguing
128
preference for secondary amines, serving as a complement to the leading precious-
metal catalysts, which work best with primary amines. In fact, the TONs of up to
600 for tertiary amide production far exceed the performance of any catalyst to date.
Perhaps most importantly, it displays notable success in the notoriously difficult
production of formamides from methanol and amines. In this manner, these results
represent a significant advancement toward a more widespread use of methanol as a
versatile and potentially renewable carbon building block in organic syntheses.
129
2.3: Iron-Catalyzed Urea Synthesis: Dehydrogenative Coupling of Methanol and
Amines
2.3.1: Introduction
Ureas are frequent intermediates in organic synthesis and appear in resin
precursors, dyes, and agrochemicals.87,88 However, their most prominent role is in
pharmaceuticals, where they are key functional groups in a host of medicinal
compounds89,90 such as antibacterial,91 antiatherosclerotic,92 antidepressant,93 and
anticancer94,95 drugs. The most prevalent metal-free methods for synthesizing ureas use
highly toxic or dangerous starting materials, such as phosgene and its derivatives,
isocyanates, azides,96-98 or carbon monoxide (CO).97,99,100 Such CO reactions also
necessitate elevated temperatures and pressures or the concurrent use of strong
oxidants.97,101-106 Carbon dioxide (CO2) can be used as a non-toxic urea precursor, but these
reactions typically require high temperatures and pressures, expensive dehydrating agents,
Scheme 2.3.1: Metal-free pathways for urea formation.
130
or strong bases.97,107-109 Additionally, all current metal-free methods produce
stoichiometric amounts of inorganic salts as waste (Scheme 2.3.1).
Transition metal catalysis is a potential strategy for improving the preparation of
ureas (Scheme 2.3.2). Of the most highly explored methods, the synthesis of symmetric
ureas via the metal-catalyzed oxidative carbonylation of amines requires high catalyst
loadings, generally gives low yields, and often forms significant amounts of side products
(such as oxamides, carbamate esters, or CO2). Furthermore, elevated pressures of CO and
harsh oxidizing conditions are needed.100,107,110-119 Urea generation through the metal-
catalyzed reaction of CO2 and substituted amines suffers from high temperatures or
pressures, base additives, poor reaction yields and scope, or some combination of these
factors.107,120-130 Metal-catalyzed dehydrogenative coupling approaches for urea synthesis
avoid pressurization concerns and have the advantage of producing dihydrogen (H2) as the
only byproduct. Unfortunately, the few examples of dehydrogenative coupling of
formamides and amines131-134 and amide cross-coupling135,136 require high catalyst loadings
Scheme 2.3.2: Metal-catalyzed urea synthesis from non-methanol sources.
131
(2-5 mol%) and most also need a stoichiometric peroxide additive. In addition, the
preparation of the formamide starting materials is often tedious and typically employs
harsh formylating reagents.139 A desirable alternative is to instead dehydrogenatively
couple methanol with an amine to form a formamide intermediate. Subsequent
dehydrogenative coupling of the formamide with another equivalent of amine generates a
urea (and H2; Scheme 2.3.3). The paucity of catalysts suitable for even the first step of this
transformation was discussed in the previous section, and is demonstrated by the fact that
while numerous metal catalysts are capable of dehydrogenative amidation,137 only four
exhibit any activity with respect to coupling methanol and amines to formamides in the
absence of harsh oxidative conditions.138-141 Only the five-coordinate iron catalyst 1 does
so with turnover numbers (TONs) greater than 50.141 The incorporation of primary amines
into dehydrogenative amidation and ureation reactions faces the additional challenge of
potential imine formation via dehydration from a plausible hemiaminal intermediate.142
While imines can be valuable and dedicated dehydrogenative imination catalysts exist,143-
145 in the case of ureation they represent undesirable side products that detrimentally affect
selectivity. As primary amines are required for dehydrogenative ureations that proceed
through isocyanate intermediates, the minimization or elimination of imination further
restricts the suite of suitable catalysts.
Scheme 2.3.3: Metal-promoted pathway for dehydrogenative urea formation from
methanol and amines.
132
Recently, Kim and Hong described the sole example of catalytic urea formation
directly from methanol and amines (Figure 2.3.1).146 Their precious metal ruthenium
catalyst exhibits high selectivities and excellent yields for the production of symmetric and
unsymmetric ureas, with TONs of up to 190 and 15, respectively, and its applicability
toward reversible147 hydrogen storage has been demonstrated.148 However, the synthesis of
unsymmetric ureas requires high catalyst loadings and a complicated two-step process. To
date, no other base-metal catalysts for the production of ureas from methanol and amines
have been described, with the most closely related examples instead relying on formamide
starting materials133 or isocyanate reagents.149 Base-metal catalysts for dehydrogenative
urea synthesis are particularly desirable for pharmaceutical applications where there are
stringent requirements regarding product separation from toxic metals.150 Previous work in
our laboratories as well as results introduced in the amidation section of this chapter have
verified the catalytic ability of iron-pincer complexes of the type (RPNP)Fe(H)(CO) (RPNP
= N[CH2CH2(PR2)]2, R = iPr, Cy) for both methanol dehydrogenation151 and the
dehydrogenative amidation of alcohols with secondary amines,141 suggesting the
applicability of this class of compounds toward ureation. In this section, the use of complex
1 ((iPrPNP)Fe(H)(CO), Figure 2.3.1) in the first base-metal-catalyzed dehydrogenative
Figure 2.3.1: Ruthenium and iron catalysts for symmetric and unsymmetric urea
formation from the dehydrogenative coupling of methanol and amines.
133
coupling of methanol and primary amines to selectively form ureas will be discussed. This
system exhibits excellent activities in the absence of any additives and has been employed
to isolate ureas on a scale of several hundred milligrams, corresponding to TONs of up to
160 for symmetric ureas and 176 for unsymmetric ureas.
2.3.2: Results and Discussion
Catalytic Studies
Due to the relative insolubility of 1,3-dicyclohexylurea in most organic solvents
and the corresponding ease of post-reaction isolation, initial investigations into the ability
of 1 to catalyze the formation of symmetric ureas employed methanol and cyclohexylamine
as the substrates. A promising TON of 23 for the formation of 1,3-dicyclohexylurea was
observed using a one to four molar ratio of methanol to amine and 0.5 mol% of 1 in 5 mL
tetrahydrofuran (THF) at 80 °C for 8 hours. Empirical optimization of the reaction resulted
in minimal improvements (Tables A2.9-A2.13) except for increasing the reaction
temperature, which nearly tripled the TON for urea formation to 66 (Table 2.3.1, entry 12).
While the total yield of urea is modest (ca 33%), sterics may play a limiting kinetic role
given the need to attach two amine substrates to the methanol-derived carbonyl carbon of
the urea.
The encouraging preliminary results prompted examination of the substrate scope
for dehydrogenative ureation (Table 2.3.1). The most productive substrates were terminal
alkylamines, with pentylamine giving the highest TON of 160 (80% yield, entry 1). Small
steric changes such as using a branched alkylamine (isobutylamine, entry 2) or elongating
the alkyl chain (heptylamine, entry 3) did not significantly alter catalyst performance.
Likewise, capping the alkylamine chain with a methoxy group (2-methoxyethylamine,
134
Entry Amine TONc Yield (%)
1 160 80%
2 150 75%
3 147 74%
4 144 72%
5 12d 6%
6b 156e 78%
7 140 70%
8 126 63%
9b 90 45%
10 123 62%
11 116 58%
12b 66 33%
13 22e,f 11%
14 0 --
15 0 --
aReaction conditions: 3 mmol methanol, 12 mmol amine, 0.5
mol% 1, 5 mL THF at 120 °C for 8 hours. Each entry is an
average of two trials unless otherwise indicated. bAverage of
three trials. cBased on yield of isolated urea of >99% purity
(as determined by 1H NMR spectroscopy) unless otherwise
indicated. d >98% purity. e >97% purity. fMixture of isomers,
~75:25 cis:trans.
Table 2.3.1: Dehydrogenative Symmetric Urea
Formation from Methanol and Amines Catalyzed by 1.a
135
entry 4) afforded good yields. However, switching from a terminal amine (entry 3) to an
internal amine (2-aminoheptane, entry 5) significantly decreased the TON. While
amidation proceeded favorably (TON = 53), it is possible that the sterically hindered
internal amine significantly increased the barrier either for attack by a second equivalent
of amine or for the dehydrogenation of intermediates necessary to form the corresponding
symmetric urea.
Electronic influences were probed using a series of benzylamine derivatives
(entries 6-10) and it was found that the presence of electron-donating substituents such as
methoxy and methyl in the para position of the benzyl moiety enhanced the TON (entries
6 and 7, respectively) compared to unsubstituted benzylamine (entry 8). An electron-
withdrawing substituent such as trifluoromethyl slightly decreased the TON (entry 9).
These substituent changes affect the nucleophilicity of the amine substrate which likely
explains this trend in TON. Attempts to decrease steric hindrance by switching from
benzylamine to 2-phenethylamine had little effect (entries 8 versus 10, respectively),
although it is possible that the increased flexibility in the short carbon chain could
counteract some of the desired steric relief. Entries 11 and 12 demonstrate that while
smaller ring structures can still perform reasonably well, there is significant steric
hindrance that increases rapidly with size, as the change of a cyclopentyl group to a
cyclohexyl group dropped the TON by nearly half. The attempted synthesis of a cyclic urea
(entry 13) from a diamine was successful, although it gave a poor yield compared to its
monoamine counterpart (entry 12). This could be partially due to the cis/trans mixture of
the starting amine (~60:40), as formation of the strained trans product is less favorable
compared to its cis counterpart (though both cis and trans products were observed), thereby
136
detracting from catalytic performance. However, analogous experiments by Kim and Hong
displayed a similar reduced activity toward diamines that was attributed to amine
coordination effects, which could also be involved here.146 Finally, aniline and its methyl
derivative (entries 14 and 15) reacted poorly under these conditions, giving very little of
the corresponding formamides and no conversion to the desired ureas. This is attributed to
a lack of nucleophilicity from the poor electron donation of the aryl substituent, which has
been a common obstacle for catalytic dehydrogenative coupling reactions using amines.152
Addition of a co-catalytic amount of exogenous base did little to overcome these limitations
(Table A2.14). Overall, 1 represents one of only two examples of metal complexes capable
of catalyzing urea synthesis from alcohols and primary amines. It has the advantage of
containing a cheaper, more abundant base metal while still displaying good yields in the
production of symmetric ureas from a range of substrates, without the formation of imines.
Furthermore, in several cases 1 affords isolated pure urea product on a synthetically useful
scale.
Encouraged by the performance of 1 in the synthesis of symmetric ureas, further
investigations addressed the preparation of unsymmetric ureas,153 which are most prevalent
as key functional groups in pharmaceuticals.93 While unsymmetric ureas can be acquired
via the transamidation of ureas, either by metal-free97 or metal-catalyzed153 means, these
methods have limited scope, can require base or reductant additives, and suffer from the
same toxicity issues in generating the urea starting materials. Kim and Hong accessed these
products from methanol and amines using a one-pot, two-step method that, while effective,
required significant and sequential catalyst loadings (6 mol% total, added in two portions)
and was restricted to benzylamine as a substrate. The direct reaction of methanol with two
137
different primary amines has been shown to result in a distribution of symmetric and
unsymmetric ureas,146 so our approach focused on selectively forming unsymmetric ureas
from the reactions of formamides and amines (Table 2.3.2). As previously mentioned, the
few metal-catalyzed examples of this process exhibit very low TONs (<50).132-134 While
starting with formamides is not ideal due to cost and ease of synthesis, the potential for
generating these materials from the dehydrogenative coupling of methanol and amines was
Entry R R’ Yieldb (%) Conv.c (%) Sel.d (%)
1 85% 89% 96%
2 86% 93% 92%
3
49% 79% 68%
4 85% 86% 82%
5
11%
28%
78%
6
0% 16% --
Table 2.3.2: Dehydrogenative Unsymmetric Urea Formation from Formamides and
Amines Catalyzed by 1.a
aReaction conditions: 3 mmol formamide, 3 mmol amine, 0.5 mol% 1, 5 mL THF at 120 °C for 16 hours.
Each entry is an average of two trials. bIsolated yield. cBased on formamide consumption. dSelectivity:
percentage unsymmetric urea (compared to symmetric ureas) in final product.
138
illustrated in the previous section of this chapter and may serve to alleviate some of the
reservations associated with this aspect of the method.
Re-optimization of the reaction conditions for producing unsymmetric rather than
symmetric ureas using 1 resulted in a change of the ratio of starting materials (to 1:1
formamide:amine) and in reaction time from 8 hours to 16 hours (Tables A2.15-A2.17).
Initial experiments involving benzylformamide and cyclohexylamine gave excellent yields
of the desired unsymmetric urea (Table 2.3.2, entry 1) with high selectivity. However, it
was found that some scrambling of the starting formamide or the unsymmetric product (or
both) had occurred to give trace amounts of both symmetric ureas (Scheme 2.3.4). Further
experiments under catalytic conditions showed that scrambling of the starting formamide
can occur in small amounts in the absence of catalyst, but suggested an iron-catalyzed
enhancement of this process (Table A2.18). NMR experiments also indicated that while
hydrogenation of the unsymmetric urea product to the opposing formamide and amine pair
could occur, it was a very minor process (ca 1.5% conversion) under the conditions studied
(10 mol% 1, THF-d8, 120 °C, 16 hours, 1 atm H2; Figure A2.3). Additional NMR
experiments revealed that the urea product can undergo further reactions with both amines
Scheme 2.3.4: Potential scrambling pathways in unsymmetric urea formation.
139
and formamides in solution, providing yet another pathway to the observed scrambling
(Figures A2.4-A2.6). While insightful, these experiments did not distinguish whether
scrambling of the starting formamide or of the product urea was the predominant process
for producing the undesirable symmetric ureas (Figure A2.7). Catalytic trials involving the
opposing reactant pair of cyclohexylformamide and benzylamine (Table 2.3.2, entry 3)
revealed a significant decrease in selectivity, presumably due to a corresponding increase
in the rate of scrambling compared to unsymmetric ureation.
Efforts to enhance yield through minor substituent manipulation did not have a
significant effect. While 4-methoxybenzylamine and pentylamine were both high-
performing substrates for the formation of symmetric ureas, there was little advantage to
the N-(4-methoxybenzyl)formamide and pentylamine combination over benzylformamide
and cyclohexylamine in either TON or selectivity (Table 2.3.2, entry 2). This indicated that
starting from combinations of formamides and amines that minimized scrambling (for
example, benzylformamide or its derivatives rather than cyclohexylformamide) was far
more important than small electronic or steric changes in substituents. As a result, other
experiments with respect to unsymmetric urea production focused on species that were
reminiscent of medicinally-relevant ureas (entries 4-6).93
Although the initial reaction of isobutylformamide and ethanolamine did not yield
urea, changing the –OH group on the amine to a methoxy group (entry 4) gave excellent
turnover, albeit with poorer selectivity than for formamides containing benzyl derivatives.
The failure of ethanolamine was attributed to its preference for forming an iron-alkoxide
species with the catalyst.141 The production of unsymmetric ureas containing an aniline
moiety on one side is highly desirable in a variety of medical applications including for
140
antitumor and anticonvulsant agents,91,94,95 however, formanilide was a poor substrate for
dehydrogenative coupling and scrambling to form the more active pentylformamide was
observed as the primary process (entry 5). A similar lack of nucleophilicity is reflected in
the amine of entry 6 as only formamide scrambling was observed with benzylformamide
(a known active substrate). Overall, while the possibility of obtaining unsymmetric ureas
in high yields with good selectivities was demonstrated, it was found that kinetic control
was far more important than qualitative changes in sterics or electronics. As a result, a full
substrate scope would not be as informative and was not performed. Regardless, 1 exhibits
TONs over three times greater than the existing catalysts for this method of obtaining
unsymmetric ureas132-134 and the one-pot two-step method starting from alcohol and
amine,146 and has the added benefits of being both base-metal-catalyzed and additive free.
The occurrence of reagent and product scrambling in the generation of unsymmetric ureas
is a drawback to this application of catalyst 1. However, this issue is not unique to iron and
further understanding of ways to control scrambling outside of reagent selection will
greatly expand the substrate scope and applicability of metal-mediated unsymmetric urea
formation.
Mechanistic Investigations
There are two main pathways proposed for urea formation, as shown in Scheme
2.3.5. Kim and Hong postulated that after initial dehydrogenative amidation (that can
proceed via either a hemiaminal or hemiacetal intermediate), their ruthenium-catalyzed
ureation followed path (a) to generate an aminal, followed by metal-catalyzed
dehydrogenation to yield the urea product.146 Alternatively, isocyanate intermediates
(usually metal-bound) have been implicated in other metal-mediated urea formation
141
reactions starting from amides,132,136 as well as in the majority of metal-free ureation
cases.96,97 These examples would suggest path (b), whereby urea generation stems from
metal-catalyzed dehydrogenation of a formamide to an isocyanate, followed by
nucleophilic attack of an amine. Mechanistic investigations were therefore performed
using 1 to elucidate whether iron-catalyzed urea formation proceeds via an aminal or an
isocyanate intermediate. Our attempts to make a tetrasubstituted urea species
(R2NC(O)NR’2) from a tertiary amide and a secondary amine were unsuccessful (Figure
A2.8, Table A2.19). This could be due to the inability of the tertiary amide to undergo
dehydrogenation to the isocyanate, although steric hindrance in generating a putative
aminal intermediate cannot be ruled out. Therefore, a similar analysis was performed using
two synthetic routes for the preparation of 1,1,3-tripentylurea (Scheme 2.3.6). While the
NMR-scale reaction of pentylformamide and dipentylamine displayed approximately 80%
conversion to the desired urea,86 the alternative combination of dipentylformamide and
pentylamine did not produce any coupling product (Figures A2.9-A2.10). The main
Scheme 2.3.5: Proposed mechanistic pathways for metal-catalyzed urea formation
starting from methanol and primary amines.
Scheme 2.3.6: Formamide-dependent synthesis of 1,1,3-tripentylurea.
142
difference between these two routes is the ability of the starting formamide to form an
isocyanate intermediate, as the sterics of the final product are identical. Further evidence
for an isocyanate intermediate was obtained through the reaction of cyclohexylformamide
with 1 in the absence of other reagents. In this reaction the iron-dihydride species
(presumably generated from formamide dehydrogenation) was observed by NMR
spectroscopy, along with cyclohexyl isocyanate (confirmed by GC analysis) as the sole
organic product (Figure A2.11). These results indicate that iron-catalyzed dehydrogenative
ureation most likely proceeds through an isocyanate intermediate as shown in Scheme
2.3.5b.
2.3.3: Summary
Complex 1 displays excellent catalytic ability for urea formation via the
dehydrogenative coupling of methanol and amines. While most precious-metal catalysts
for the coupling of alcohols and primary amines stop at the formamide, this base-metal
species has a rare propensity to undergo further coupling to the urea. The performance of
1 in the synthesis of symmetric ureas from methanol and primary amines rivals that of the
only other (ruthenium catalyzed) example, and its TONs in the synthesis of unsymmetric
ureas from formamides and primary amines are several times better than those of any other
existing catalysts for that synthetic route. In addition, 1 performs these dehydrogenative
reactions without the use of an exogenous base or a hydrogen acceptor, enhancing its
environmental and economic profile. In contrast to the proposed behavior of the ruthenium
catalyst, mechanistic studies suggest that in this case the final ureation step proceeds via
an isocyanate intermediate. Trapping the transient isocyanate species with other reagents
(such as alcohols to give urethanes)154 may provide alternative synthetic pathways to a suite
143
of desirable compounds starting from dehydrogenative coupling. Despite limitations with
respect to scrambling and steric sensitivity, this works represents yet another example of
the superiority of 1 with respect to methanol activation and subsequent transformation into
high-value products.
144
2.4: Chapter Summary
In this chapter, the superior ability of the iron pincer complex (iPrPNP)Fe(H)(CO)
(complex 1) in catalytic dehydrogenative coupling reactions for organic synthesis was
established. Complex 1 overcame the inherent energetic challenges to methanol activation
that thus far have limited its more widespread incorporation as a C1 carbon source in
dehydrogenative organic transformations. As a result, 1 emerges as the most active base-
metal catalyst for dehydrogenative amidation, with TONs that exceed its precious metal
competition with respect to secondary amines. In addition, 1 is the first base-metal catalyst
for dehydrogenative urea formation from methanol and primary amines and represents one
of only two catalysts for this method of urea synthesis, both exhibiting comparable TONs.
No base or hydrogen acceptor additives are needed in either amidation or ureation, making
this method of methanol utilization applicable to a wider range of functionalized species.
The significance of these results rests in the pivotal role of formamide and urea
functionalities in medicinal applications and the potential for sourcing such valuable
chemicals from a cheap and abundant carbon source in a more environmentally benign
manner using a non-precious-metal catalyst. That being said, understanding and solving
the severe steric limitations of catalyst 1 (possibly through modifications of the ancillary
ligand environment), improving its reactivity with less nucleophilic substrates, and
developing methods of increasing product selectivity in ureation are key issues that must
be addressed before its more widespread use in formamide and urea production is possible.
Additionally, while the performance of catalyst 1 compares with and often exceeds its
precious-metal competition for select substrates, the activities of base-metal catalysts in
general in this field of application are an order of magnitude worse than those of the best
145
precious-metal catalysts. Ways of circumventing this problem (such as increased
recyclability of base-metal catalysts) will need to be established before implementation of
these methods on a large scale can be realized.
146
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Appendix 1
1.2.A.i.: Nickel-Promoted CO2-Alkene Coupling for the Stereoselective Production of
Substituted Cyclic Anhydrides
General Considerations
All manipulations were carried out under a nitrogen or argon atmosphere using standard
Schlenk, vacuum, cannula, or glovebox techniques. Ligand 9 and NaBAr4F were prepared
according to previously published procedures.1,2 Bulk solvents were dried and
deoxygenated using literature procedures.3 All chemicals were purchased from VWR,
Aldrich, Fisher Scientific, Strem, or Cambridge Isotope Laboratories. NMR solvents were
dried over 3 Å molecular sieves and then used without further manipulation, or sodium and
then vacuum transferred prior to use. 1H, 13C, and 31P NMR spectra were recorded on
Bruker DRX 400 or Avance 600 MHz spectrometers at ambient temperature, unless
otherwise noted. Chemical shifts are reported in ppm; J values are given in Hz. In some
cases, the inequivalent 31P nuclei have coupling constants too small to be observed due to
the combination of 2J and 3J couplings and are therefore denoted as singlets. 1H and 13C
chemical shifts are referenced to residual solvent signals; 31P NMR chemical shifts are
referenced to an external standard of H3PO4. Probe temperatures were calibrated using
ethylene glycol and methanol as previously described.4 GCMS analyses were performed
using an Agilent gas chromatogram with mass spectrometer/Agilent CP-chiraSil-DEX
chiral gas chromatography column.
157
Figure A1.1: Structure key for ligand numbering. iPr = isopropyl. tBu = tert-butyl.
1 2
3 4 5
6 7
8 9
158
Ligand Number Name
1 (R)-2,2’-bis[bis(4-methoxy-3,5-di-tert-
butyl-phenyl)phosphino]-4,4’,6,6’-
tetramethoxybiphenyl =
(R)-DTBM-GARPHOS
2 1,2-bis((2S,5S)-2,5-
dimethylphospholano)ethane =
(S,S)-Me-BPE
3 1,2-bis((2S,5S)-2,5-
diisopropylphospholano)ethane =
(S,S)-iPr-BPE
4 (-)-1,2-bis((2S,5S)-2,5-
diisopropylphospholano)benzene =
(S,S)-iPr-DUPHOS
5 (1R,1’R,2S,2’S)-2,2’-di-tert-butyl-
2,3,2’,3’-tetrahydro-1H,1’H-
(1,1’)biisophosphindolyl =
(1R,1’R,2S,2’S)-DuanPhos
6 (S,S)-1,2-bis(tert-
butylmethylphosphino)ethane =
(S,S)-tBu-BISP*
7 (R)-1-[(Sp)-2-
(dicyclohexylphosphino)ferrocenyl]
ethyldi-tert-butylphosphine =
(R)-Fc(PCy2)(PtBu2)
8 (R)-1-[(Sp)-2-
(dicyclohexylphosphino)ferrocenyl]
ethyldicyclohexylphosphine =
(R)-Fc(PCy2)2
9 1,1’-bis((2S,5S)-2,5-
diisopropylphospholano)ferrocene =
(S,S)-Fc-iPr-BPE
Synthesis of Ligand-Ni(COD) Species and Select NMR Data
Scheme A1.1: Synthesis of Ligand-Ni(COD) species.
Table A1.1: Naming Key for Ligand Numbering.
159
Synthesis of 1-Ni(COD):
A J. Young NMR tube was treated with 1 (0.010 g, 0.008 mmol), Ni(COD)2 (0.003 g, 0.011
mmol), and C6D6 (0.5 mL). The solution was heated overnight at 60 ̊ C and then filtered
through Celite. 1H NMR (400 MHz, C6D6): 1.44 (d, J = 14 Hz), 2.90 (s), 3.41 (t, J = 8),
6.71 (s), 7.56 (s), 7.73 (s). 31P{1H} NMR (400 MHz, C6D6): 36.0 (s).
Synthesis of 2-Ni(COD):
A 20 mL scintillation vial was charged with Ni(COD)2 (0.020 g, 0.073 mmol) and C6D6,
with stirring. The solution was transferred to a J. Young NMR tube and then 1 equivalent
of 2 (20 μL, 0.073 mmol) was injected using a microsyringe. 1H NMR (400 MHz, C6D6):
0.58 (m), 0.80 (m), 1.31 (m), 1.50 (m), 1.73 (m), 2.40 (m), 2.72 (m), 4.04 (m, 2-Ni(COD)),
5.06 (m, 2-Ni(COD)). 31P{1H} NMR (300 MHz, C6D6): 68.7 (s).
Synthesis of 3-Ni(COD):
A 20 mL scintillation vial was charged with Ni(COD)2 (0.051 g, 0.19 mmol) and THF (3
mL). A separate scintillation vial was charged with 3 (0.069 g, 0.19 mmol) and THF (3
mL). The ligand solution was added dropwise with stirring to the Ni(COD)2 solution, then
stirred overnight at room temperature. The THF was removed under vacuum and the
resulting oily yellow solid was dissolved in pentane and cooled to -35 ̊ C. The solution was
filtered through Celite and the pentane was removed under vacuum to yield a yellow solid.
1H NMR (400 MHz, C6D6): 0.74 (d, J = 6 Hz, CH(CH3)2), 0.84 (m), 0.96 (m), 1.02 (m),
1.26 (m), 2.39 (m), 2.73 (m), 4.13 (m, 3-Ni(COD)), 4.92 (m, 3-Ni(COD)). 31P{1H} NMR
(400 MHz, C6D6): 58.6 (s).
160
Synthesis of 4-Ni(COD):
A 20 mL scintillation vial was charged with Ni(COD)2 (0.033 g, 0.12 mmol) and THF (3
mL). A separate scintillation vial was charged with 4 (0.050 g, 0.12 mmol) and THF (3
mL). The ligand solution was added dropwise, with stirring, to the Ni(COD)2 solution, then
transferred to a 50 mL round bottom flask fitted with a needle valve. The solution was
heated at 50 ̊ C overnight and then the solvent was removed under vacuum. The resulting
orange oil was dissolved in pentane, cooled to -35 ̊ C, and filtered through Celite. The
pentane was removed under vacuum to give an oily orange solid. 1H NMR (400 MHz,
C6D6): 0.52 (s), 0.65 (s), 0.83 (s), 1,10 (m), 1.22 (m), 1.36 (m), 1.50 (m), 1.62 (m), 2.43
(m), 2.74 (s), 5.06 (s), 7.56 (s). 31P{1H} NMR (400 MHz, C6D6): 56.3 (s).
Synthesis of 5-Ni(COD):
A 20 mL scintillation vial was charged with Ni(COD)2 (0.028 g, 0.10 mmol) and THF (3
mL). A separate scintillation vial was charged with 5 (0.039 g, 0.10 mmol) and THF (3
mL). The ligand solution was added dropwise with stirring to the Ni(COD)2 solution, then
stirred overnight at room temperature. The THF was removed under vacuum and the
resulting oily yellow solid was dissolved in pentane and cooled to -35 ̊ C. The solution was
filtered through Celite and the pentane was removed under vacuum to yield a yellow solid.
1H NMR (400 MHz, C6D6): 0.82 (d, J = 9 Hz, C(CH3)4), 2.27 (s), 2.80 (s, 5-Ni(COD)),
3.08 (s), 3.95 (m), 4.57 (s, 5-Ni(COD)), 7.05 (s, 4H, aryl), 7.23 (s, 4H, aryl). 31P{1H} NMR
(400 MHz, C6D6): 95.1 (s).
161
Synthesis of 6-Ni(COD)
A 20 mL scintillation vial was charged with Ni(COD)2 (0.089 g, 0.32 mmol) and THF (3
mL). A separate scintillation vial was charged with 6 (0.076 g, 0.32 mmol) and THF (3
mL). The ligand solution was added dropwise with stirring to the Ni(COD)2 solution, then
stirred for four hours at room temperature. The THF was removed under vacuum and the
resulting solid was extracted with cold diethyl ether. The extractions were collected and
the ether was removed under vacuum to yield a brown oil. 1H NMR (400 MHz, C6D6): 0.94
(d, J = 5), 1.04 (d, J = 12), 1.20 (m), 1.39 (m), 2.17 (m), 2.42 (m), 2.68 (m), 4.07 (m), 4.20
(m), 4.48 (m), 4.55 (m). 31P{1H} NMR (400 MHz, C6D6): 51.9 (minor, s), 53.1 (major, s).
Synthesis of 7-Ni(COD)
Due to lack of reaction at room temperature and poor conversion followed by
decomposition at elevated temperatures (60-80 °C), this species was not isolated and was
instead formed in situ during coupling reactions. 31P{1H} NMR (400 MHz, THF/C6D6):
21.8 (d, J = 27 Hz, 1P), 80.2 (d, J = 27 Hz, 1P).
Synthesis of 8-Ni(COD)
A 20 mL scintillation vial was charged with Ni(COD)2 (0.011 g, 0.04 mmol) and THF (3
mL). A separate scintillation vial was charged with 8 (0.025 g, 0.04 mmol) and THF (3
mL). The ligand solution was added dropwise with stirring to the Ni(COD)2 solution, then
stirred overnight at room temperature. The THF was removed under vacuum and the
resulting oily solid was washed with pentane and re-dried to yield a red-orange powder. 1H
NMR (400 MHz, C6D6): 1.03 (m), 1.28 (m), 1.52 (m), 1.64 (m), 1.75 (m), 2.54 (m), 2.88
(m), 4.02 (m), 4.09 (s), 4.10 (d, J = 6), 4.43 (s), 4.62 (s), 5.08 (br s), 5.17 (br s). 31P{1H}
162
NMR (400 MHz, C6D6): 25.1 (minor a, d, J = 28, 1P), 25.7 (minor b, d, J = 29, 1P), 26.1
(major, d, J = 24, 1P), 48.0 (minor b, d, J = 29, 1P), 49.0 (major, d, J = 24, 1P), 49.1 (minor
a, d, J = 28, 1P).
Synthesis of 9-Ni(COD)
A 20 mL scintillation vial was charged with Ni(COD)2 (0.012 g, 0.04 mmol) and THF (3
mL). A separate scintillation vial was charged with 9 (0.023 g, 0.04 mmol) and THF (3
mL). The ligand solution was added dropwise with stirring to the Ni(COD)2 solution, then
stirred overnight at room temperature. The THF was removed under vacuum yielding a
yellow powder. 1H NMR (400 MHz, C6D6): 0.75 (s), 0.99 (s), 1.05 (s), 1.30 (s), 1.35 (s),
1.41 (s), 1.57 (s), 1.68 (s), 1.79 (s), 2.01 (s), 2.28 (s), 2.40 (s), 2.51 (s), 3.07 (br s), 3.88 (s),
4.03 (s), 4.18 (s), 4.56 (s), 4.60 (s), 4.71 (s). 31P{1H} NMR (400 MHz, C6D6): 26.2 (major,
d, J = 32, 1P), 26.9 (minor, d, J = 34, 1P), 28.4 (s, overlapping with minor d), 29.2 (major,
d, J = 32, 1P).
Synthesis of Ligand-Ni-ethylene and Ligand-Ni-lactone Species
*Note: the table after the general methods sections gives the optimized reaction conditions
for each specific ligand
Scheme A1.2: Synthesis of Ligand-Ni-lactone species.
163
General Method: Ligands 1-6 and 8-9
Dissolve 5-10 mg of Ligand-Ni(COD) and 1 equiv. NaBAr4F (if needed) in C6D6, acetone-
d6, or THF-d8 and transfer to a J. Young NMR tube (0-4 atm intended total pressure) or a
high-pressure NMR tube (5-13 atm intended total pressure). Sequentially add C2H4, then
CO2 or 13CO2. Heat up to 80 °C as necessary for 1-5 days.
Adaptation for ligand 7: Instead of the Ligand-Ni(COD) species, combine 5 mg Ni(COD)2
and 10 mg ligand with 1 equiv. NaBAr4F in the chosen NMR solvent.
Synthesis of Ligand-Ni-propylene and Ligand-Ni-lactone’ Species
General Method:
Dissolve 5-10 mg of Ligand-Ni(COD) and 1 equiv. NaBAr4F (if needed) in C6D6, acetone-
d6, or THF-d8 and transfer to a J. Young NMR tube (0-4 atm intended total pressure) or a
high-pressure NMR tube (5-13 atm intended total pressure). Sequentially add C3H6, then
CO2 or 13CO2. Heat up to 80 °C as necessary for 1-5 days.
Scheme A1.3: Synthesis of Ligand-Ni-lactone’ species.
164
Ligand LNi(COD) NaBAr4F C2H4 CO2 Solvent Temp. Time
1 - - - - - - -
2 5 mg - 11 atm - C6D6 22 °C 2 hrs.
3 10 mg 18 mg 0.1 atm 0.1 atm Acetone-d6 22 °C 12 hrs.
4 10 mg 15 mg 2 atm 2 atm THF-d8 50 °C 12 hrs.
5 7 mg 11 mg 1.5 atm 1.5 atm C6D6 22 °C 12 hrs.
6 8 mg - 5 atm 5 atm C6D6 22 °C 48 hrs.
7 5 mg
Ni(COD)2
10 mg
ligand
16 mg 2 atm 2 atm THF-d8 22 °C 48 hrs.
8 10 mg 12 mg 5 atm 5 atm THF-d8 22 °C 48 hrs.
9 - - - - - - -
Ligand LNi(COD) NaBAr4F C3H6 CO2 Solvent Temp. Time
1 - - - - - - -
2 - - - - - - -
3 15 mg - 5 atm 5 atm C6D6 22 °C 48 hrs.
4 5 mg 8 mg 8 atm 2 atm THF-d8 22 °C 48 hrs.
5 10 mg 17 mg 5 atm 5 atm THF-d8 22 °C 48 hrs.
6 10 mg 22 mg 5 atm 5 atm THF-d8 22 °C 72 hrs.
7 - - - - - - -
8 - - - - - - -
9 12 mg 18 mg 5 atm 5 atm THF-d8 50 °C 12 hrs.
Ligand Ni-ethylene
(ppm)
Ni-lactone
(ppm)
Ni-propylene
(ppm)
Ni-lactone’
(ppm)
1 - - - -
2
(C6D6)
67.5
(s, 2P)
(121 MHz)
- - -
3
(acetone-d6)
55.8
(s, 2P)
(121 MHz)
58.1
(s, 1P)
64.8
(s, 1P)
(300 MHz)
53.7
(d, J = 69, 1P)
55.4
(d, J = 70, 1P)
major 49.1
56.5
(d, J = 13, 1P),
58.8
(d, J = 13, 1P)
60.1
Table A1.2: Optimized Experimental Conditions for Ni-lactone Syntheses
Table A1.3: Optimized Experimental Conditions for Ni-lactone’ Syntheses
Table A1.4: 31P{1H} NMR Shifts for Synthesized Ni Species. J Values are in Hz.
165
(d, J = 66, 1P)
54.8
(d, J = 67, 1P)
minor (162 MHz)
(d, J = 12, 1P),
64.8
(d, J = 12, 1P)
60.9
(d, J = 3, 1P)
62.1
(d, J = 3, 1P)
(162 MHz)
4
(THF-d8)
60.2
(s, 2P)
(243 MHz)
59.0
(s, 1P)
63.4
(s, 1P)
(243 MHz)
54.8
(d, J = 69, 1P),
56.6
(d, J = 69, 1P)
(162 Mz)
-
5
(C6D6, C2H4)
(THF-d8, C3H6)
97.4
(s, 2P)
(162 MHz)
91.3 (d, J = 10,
1P)
105.7 (d, J = 9,
1P)
(162 MHz)
Possible: 94.7
(s, 2P)
(162 MHz)
-
6
(C6D6, C2H4)
(THF-d8, C3H6)
53.8
(s, 2P)
(162 MHz)
50.8 (s, 1P)
59.1 (s, 1P)
(162 MHz)
60.6 (s, 1P)
62.8 (s, 1P)
(162 MHz)
51.1 (d, J = 13,
1P))
64.3 (s, 1P)
(162 MHz)
7
(THF-d8)
26.6 (d, J = 24,
1P)
82.9 (d, J = 25,
1P)
(162 MHz)
- - -
8
(THF-d8)
28.0 (d, J = 21,
1P)
52.1 (d, J = 22,
1P)
(162 MHz)
29.8 (d, J = 24,
1P)
54.0 (d, J = 24,
1P)
(162 MHz)
- -
9
(THF-d8)
- - 31.7 (d, J = 29,
1P),
32.6 (d, J = 29,
1P)
major
27.6 (d, J = 29,
1P),
32.1 (d, J = 28,
1P)
minor
(162 MHz)
-
166
Release of Cyclic Anhydride Product from Ni-lactone and Ni-lactone’ Species
General Method:
Take the Ligand-Ni-lactone and Ligand-Ni-lactone’ solutions and treat with 1 atm of CO
gas in a J. Young or high-pressure NMR tube. Allow to react for approximately 20 minutes
at room temperature. Analyze for SA or MSA using 1H and 13C NMR spectroscopy and
GCMS analysis.
Ligand Ligand-Ni(CO)2 (ppm) Solvent/Spectrometer
1 - -
2 - -
3 55.5 acetone-d6/400 MHz
4 58.9 acetone-d6/400 MHz
5 96.2 acetone-d6/400 MHz
6 53.2 acetone-d6/400 MHz
7 22.7
(d, J = 8)
88.3
(d, J = 8)
THF-d8/400 MHz
8 29.8
(d, J = 11)
52.7
(d, J = 11)
acetone-d6/400 MHz
9 28.6
41.4
THF-d8/600 MHz
Scheme A1.4: Release of cyclic anhydride product from lactone using CO gas.
Table A1.5: 31P{1H} NMR Shifts for Ligand-Ni(CO)2 Species. J Values are in Hz.
167
Peaks/GC Information Used in the Identification of Succinic Anhydride and
Methylsuccinic Anhydride:
Succinic anhydride: 1H NMR (400 MHz, acetone-d6): 3.05 (s).
Methylsuccinic anhydride: 13C NMR (400 MHz, C6D6): 169.82 (C=O), 174.34 (C=O).
GCMS in acetone-d6: peaks at 16.577 and 16.866 minutes on chiral column (initial temp.:
40 ̊ C, ramp rate: 5 ̊ C/min, max. temp.: 150 ̊ C); m/z: 15, 28, 42, 55; (m+1)/z: 71.
168
Select NMR Data: L-Ni(COD) Reactions with C2H4 or C3H6, CO2, and CO
*See Tables A1.2 and A1.3 for experimental conditions; see Schemes A1.2-A1.4 for
generic reaction depictions
Figure A1.2: 31P{1H} NMR spectrum in C6D6 showing limited conversion to 1-Ni(COD). 300 MHz
spectrometer.
Figure A1.3: 31P{1H} NMR spectrum in C6D6 showing limited conversion to 2-Ni-ethylene under 11
atm of C2H4. 300 MHz spectrometer.
2-Ni(COD)
2-Ni-2 2-Ni-ethylene
1-Ni(COD)
Ligand
Impurity
Free
Ligand
169
Figure A1.4: 31P{1H} NMR spectra in acetone-d6 showing conversion from the 3-Ni(COD) species (A,
blue) to 3-Ni-ethylene (B, red), 3-Ni-lactone (C, green), and 3-Ni(CO)2 (D, purple). Blue = 600 MHz
spectrometer, others = 300 MHz.
3-Ni(COD) + NaBAr4F
+ C2H4
+ CO2
+ CO D
B
A
C C
Figure A1.5: 31P{1H} NMR spectra in acetone-d6 showing conversion from the 3-Ni(COD) species (A,
blue) to 3-Ni-propylene (B/B’, red), 3-Ni-lactone’ (C/C’/C”, green), and 3-Ni(CO)2 (D, purple). 400
MHz spectrometer.
A
D
B B’ B’
B
C” C” C’ C’ C C
3-Ni(COD) + NaBAr4F
+ C3H6
+ CO2
+ CO
170
Figure A1.6: 31P{1H} NMR spectra in THF-d8 showing conversion from the 4-Ni(COD) species (A,
blue) to 4-Ni-ethylene (B, red), 4-Ni-lactone (C, green), and 4-Ni(CO)2 (D, purple). 600 MHz
spectrometer. * = unidentified side product.
A
B
C C
D
* *
4-Ni(COD) + NaBAr4F
+ C2H4
+ CO2
+ CO
Figure A1.7: 31P{1H} NMR spectra in THF-d8 showing conversion from the 4-Ni(COD) species (A,
blue) to 4-Ni-propylene (B, blue), and 4-Ni-CO2 adduct (C, red). 400 MHz spectrometer. * = unidentified
side product. Corresponding 4-Ni(CO)2 complex is the same as (D) in the purple spectrum in Figure
A1.6.
4-Ni(COD) + NaBAr4F + C3H6
*
A B
C
+ CO2
C
B
171
Figure A1.8: 31P{1H} NMR spectra in C6D6 showing conversion from the 5-Ni(COD) species (A, blue)
to 5-Ni-ethylene (B, red), 5-Ni-lactone (C, green), and 5-Ni-(CO)2 (D, purple). 400 MHz spectrometer.
Purple spectrum in acetone-d6. NaBAr4F added after CO2 to enhance conversion.
5-Ni(COD)
+ C2H4
+ CO2 + NaBAr4F
+ CO
C C
B
A
5-Ni-5
Figure A1.9: 31P{1H} NMR spectra in THF-d8 showing conversion from the 5-Ni(COD) species (A,
blue) to potential 5-Ni-propylene (B, red), 5-Ni-CO2 adduct (C, green), and 5-Ni-(CO)2 (D, purple). 400
MHz spectrometer.
5-Ni(COD) + NaBAr4F
+ C3H6
+ 13CO2
+ CO
A
B
D
C C
D
172
Figure A1.10: 31P{1H} NMR spectra in C6D6 showing conversion from the 6-Ni(COD) species (A, blue,
two diastereomers) to 6-Ni-ethylene (B, red), 6-Ni-lactone (C, green), and 6-Ni-(CO)2 (D, purple). 400
MHz spectrometer. Purple spectrum in acetone-d6. * = unidentified decomposition product.
6-Ni(COD)
+ C2H4
+ CO2
+ CO
A
A
B
C C
D
Figure A1.11: 31P{1H} NMR spectra in THF-d8 showing conversion from the 6-Ni(COD) species (A,
blue, two diastereomers) to 6-Ni-propylene (B, red), 6-Ni-lactone’ (C, red), and 6-Ni-(CO)2 (D, green).
400 MHz spectrometer. Green spectrum in acetone-d6.
6-Ni(COD) + NaBAr4F A
A
C
C
B B
D
+ C3H6 + 13CO2
+ CO
*
173
Figure A1.12: 31P{1H} NMR spectra in THF-d8 showing conversion from the 7-Ni(COD) species (A,
blue) to 7-Ni-ethylene (B, red), no change upon CO2 addition (green), then to 7-Ni(CO)2 (D, purple). 400
MHz spectrometer. Blue spectrum in a mixture of THF/C6D6.
7-Ni(COD)
+ NaBAr4F + C2H4
+ CO2
+ CO
A A
B B
B
D D
B
Figure A1.13: 31P{1H} NMR spectra in THF-d8 showing conversion from the 8-Ni(COD) species (A,
blue) to 8-Ni-ethylene (B, red), 8-Ni-lactone (green), and 8-Ni(CO)2 (purple). 400 MHz spectrometer.
Purple spectrum in acetone-d6.
8-Ni(COD) + NaBAr4F
+ C2H4
+ CO2
+ CO
A A
B B
C C
D D
174
Figure A1.14: 31P{1H} NMR spectra in THF-d8 showing conversion from the 9-Ni(COD) species (A,
blue) to 9-Ni-propylene (B/B’, red), and 9-Ni(CO)2 (green). Green = 600 MHz spectrometer, others =
400 MHz. * = unidentified impurity.
9-Ni(COD) + NaBAr4F
+ C3H6 + CO2
+ CO
D D
B’ B’
B B
A A
*
175
Retrosynthesis: Ligand-Ni(COD) + Succinic Anhydride or Methylsuccinic Anhydride
Retrosyntheses of Ligand-Ni-lactone or Ligand-Ni–lactone’ species using SA or MSA
followed the same general procedure:
A J. Young NMR tube was charged with Ligand-Ni(COD) (0.010 g), succinic anhydride
or methylsuccinic anhydride (1.5 equiv.), and C6D6 or acetone-d6 (0.5 mL). The NMR tube
was heated at either 60 ̊ C or 70 ̊ C for 12 hours.
Retrosynthesis: Ligand-Ni(COD) + Acrylic Acid:
Retrosyntheses of Ligand-Nilactone species using acrylic acid followed the same general
procedure:
A J. Young NMR tube was charged with Ligand-Ni(COD) (0.010 g), acrylic acid (1 equiv.
or excess), and C6D6 or acetone-d6 (0.5 mL). The reaction solution was monitored
periodically at room temperature for up to 10 days.
176
Select NMR Data: Reactions of 3-, 4-, and 5-Ni(COD) without Lewis Acid;
Retrosynthetic Reactions
Figure A1.15: 31P{1H} NMR spectra of the reaction of 3-Ni(COD) with C2H4 and CO2 in the absence
of Na salt. B = 3-Ni-ethylene, C = 3-Ni-lactone. C6D6, 300 MHz spectrometer.
3-Ni(COD) + C2H4
+ CO2
B
B
C C
Figure A1.16: 31P{1H} NMR spectrum of the reaction of 4-Ni(COD) with C2H4 and CO2 in the absence
of Na salt. A = 4-Ni(COD), B = 4-Ni-ethylene, C = 6-Ni-lactone. Acetone-d6, 300 MHz spectrometer.
B C C
A
4-Ni(COD) + C2H4 + CO2
177
Figure A1.17: 31P{1H} NMR spectrum of the reaction of 5-Ni(COD) with C2H4 in the absence of Na
salt. A = 5-Ni(COD), B = 5-Ni-ethylene. C6D6, 400 MHz spectrometer.
A B
5-Ni-5
5-Ni(COD) + C2H4
Figure A1.18: 31P{1H} NMR spectra of the reaction of 3-Ni(COD) with C2H4 and CO2 (blue) or SA
(red) to generate a nickelalactone. C = 3-Ni-lactone (γ), C’ = suspected β-lactone, D = 3-Ni(CO)2.
Acetone-d6, 400 MHz spectrometer. Slight shift changes attributed to presence of Na salt.
3-Ni(COD) + NaBAr4F + C2H4 + CO2
3-Ni(COD) + SA
C C
C’ C’
C C
D
178
Figure A1.19: 31P{1H} NMR spectra of the reaction of 4-Ni(COD) with C2H4 and CO2 (blue) or SA
(red) to generate a nickelalactone. A = 4-Ni(COD), B = 4-Ni-ethylene, C = 4-Ni-lactone, D = 4-Ni(CO)2.
Acetone-d6, 300 MHz spectrometer. * = unidentified impurity.
A
A
B
D
C C
C C *
4-Ni(COD) + C2H4 + CO2
4-Ni(COD) + SA
Figure A1.20: 31P{1H} NMR spectra of the reaction of 5-Ni(COD) with C2H4 and CO2 (blue) or SA
(red) to generate a nickelalactone. A = 5-Ni(COD), B = 5-Ni-ethylene, C = 5-Ni-lactone, D = 5-Ni(CO)2.
C6D6, 400 MHz spectrometer.
5-Ni(COD) + C2H4 + CO2
5-Ni(COD) + SA
C C
C C
D
A
B
179
Figure A1.21: 31P{1H} NMR spectrum of the reaction of 3-Ni(COD) with acrylic acid to yield multiple
unidentified products inconsistent with lactone. Acetone-d6, 300 MHz spectrometer.
Figure A1.22: 31P{1H} NMR spectra of the reaction of 3-Ni(COD) with C3H6 and 13CO2 (blue) or MSA
(red, no Na salt; green with Na salt) to generate a nickelalactone. A = 3-Ni(COD), B = 3-Ni-propylene,
C/C’/C” = 3-Ni-lactone’, D = 3-Ni(CO)2. Coupling reaction in C6D6, MSA reactions in acetone-d6; 400
MHz spectrometer. Experiment not repeated with other Ligand-Ni(COD) species because of lack of
peak agreement between methods.
A
D
B B C
C’ C’ C” C”
3-Ni(COD) + NaBAr4F + C3H6 + CO2
3-Ni(COD) + NaBAr4F + MSA
3-Ni(COD) + MSA
C
A D
180
1.2.A.ii.: The Use of CO Surrogates in Nickel-Promoted CO2-Alkene Coupling for the
Catalytic Production of Cyclic Anhydrides
General Considerations
All manipulations were carried out under a nitrogen or argon atmosphere using standard
Schlenk, vacuum, cannula, or glovebox techniques. The following species were prepared
by adaptations of previously reported procedures: dcpe-Ni(COD),5,6 dcpe-Ni-lactone’,5,6
and dcpe-Ni(CO2).7,8 Bulk solvents were dried and deoxygenated using literature
procedures.3 All chemicals were purchased from VWR, Aldrich, Fisher Scientific, Strem,
or Cambridge Isotope Laboratories. NMR solvents were dried over 3 Å molecular sieves
and then used without further manipulation, or sodium and then vacuum transferred prior
to use. 1H, 13C, and 31P NMR spectra were recorded on Bruker 300 MHz Avance II+, 300
MHz DRX, 500 MHz DRX or 600 MHz spectrometers at ambient temperature, unless
otherwise noted. Chemical shifts are reported in ppm; J values are given in Hz. In some
cases, the inequivalent 31P nuclei have coupling constants too small to be observed due to
the combination of 2J and 3J couplings and are therefore denoted as singlets. 1H and 13C
chemical shifts are referenced to residual solvent signals; 31P NMR chemical shifts are
referenced to an external standard of H3PO4. Probe temperatures were calibrated using
ethylene glycol and methanol as previously described.4 GCMS analyses were performed
using an Agilent gas chromatogram with mass spectrometer.
Synthesis of dcpe-Ni(COD):
In a 20-mL scintillation vial, dissolve 0.020 g Ni(COD)2 in ~2 mL THF. In a separate vial,
dissolve 1 equiv. dcpe ligand (0.030 g) in ~4 mL THF. Add the ligand solution to the nickel
solution dropwise, with stirring. Stir at room temperature for ~4 hours, then remove the
181
solvent under vacuum. Recrystallize the resulting oily solid from pentane to yield a
powdery yellow solid. 1H NMR (C6D6, 400 MHz): 1.21-1.33 (m), 1.66 (s), 1.71 (s), 1.73
(s), 1.82-1.85 (m), 2.52 (m), 2.68 (m), 4.46 (s). 31P{1H} NMR (C6D6, 400 MHz): 59.0 (s,
2P).
Synthesis of dcpe-Ni-lactone’:
Dissolve 0.100 g dcpe-Ni(COD) in ~6 mL THF and transfer to a 25-mL heavy-walled glass
reaction vessel. Add 1.5 atm each propylene, then CO2. Stir at 50 °C overnight. De-gas the
solution and filter through Celite. Remove the solvent under vacuum to give an oily solid.
Wash the solid with cold toluene, then cold pentane (1-2 times) until the pentane wash is
colorless. Dry the resulting orange solid under vacuum. 1H NMR (C6D6, 500 MHz): 0.91-
1.01 (m), 1.10 (s), 1.16-1.40 (m), 1.55 (s), 1.65-1.69 (m), 1.82 (m), 2.02 (m), 2.12 (m), 2.34
(s), 2.52 (s), 2.68 (s), 2.82 (s). 31P{1H} NMR (C6D6, 400 MHz): 61.2 (s, 1P), 68.4 (s, 1P).
182
Entry Ni C3H6, CO2 Temp./Time
1 7 mg 1.5 atm each 70 °C, 10 min.
2 10 mg 1.5 atm each 70 °C, 3 days
3 6 mg 1.5 atm each 50 °C, 1 hr.
60 °C, overnight
70 °C, overnight
4 7 mg 1.5 atm each 50 °C, overnight
60 °C, 2 days
70 °C, overnight
5 8 mg None 50 °C, overnight
70 °C, overnight
6 10 mg 0.5 atm each 22 °C, 3 hrs.
60 °C, 5 hrs.
70 °C, overnight
7 10 mg None 22 °C, 2-3 hrs.
8 8 mg 1 atm each 60 °C, 1 hr.
9 10 mg 1 atm each 22 °C, 1 hr.
10 10 mg 1 atm each 50 °C, 1 hr.
11 10 mg 1 atm each 50 °C, 5 hrs.
12 10 mg 0.5 atm each 60 °C, 2 hrs.
13 10 mg 0.5 atm each 60 °C, 2 hrs.
14 10 mg 0.5 atm each 22 °C, 2 hrs.
15 10 mg 1 atm each 22 °C, 6 hrs.
50 °C, 2 hrs.
General Procedure for CO Surrogate Reactions:
NMR-scale: Dissolve dcpe-Ni(COD) or dcpe-Ni-lactone’ in ~0.5 mL of the desired NMR
solvent with X equiv. [CO]. Transfer to a J. Young NMR tube (total gas pressure <4 atm)
Table A1.6: Experimental Conditions for Table 1.2.A.1 Entries
Scheme A1.5: Nickel-promoted production of cyclic anhydrides using CO surrogates
183
or a high-pressure NMR tube (total gas pressure 4-13 atm). Add Y equiv. [B]. Quickly
freeze at -196 °C and add the desired pressure of propylene, then CO2 using a calibrated
gas bulb. Thaw and then heat as needed for the duration of the reaction.
Reaction-scale: Dissolve dcpe-Ni(COD) or dcpe-Ni-lactone’ in 2-6 mL’s C6H6 in a 20-
mL scintillation vial. Add X equiv. [CO], then Y equiv. [B]. If no gases pressure or heat is
needed, stir in the vial at room temperature for the duration of the reaction. Otherwise,
transfer the reactant mixture to a 25-mL heavy-walled glass reaction vessel, quickly freeze
at -196 °C and add the desired pressure of propylene, then CO2 using a calibrated gas bulb.
Thaw and then heat as needed for the duration of the reaction.
Optimization Tables for CO Surrogate Reactions:
Entry Ni Species [CO] [CO]
equiv.
[B] [B]
equiv.
Solvent MSA
(%)
1 A NFS 3 NEt3 4.5 THF-d8 0
2 A NFS 3 NEt3 4.5 C6D6 0
3b A NFS 3 - - C6D6 0
4c B NFS 3 NEt3 4 C6D6 6.30
Table A1.7: Determination of Starting Ni Speciesa
aReaction conditions: 8-17 mg Ni species, 3 equiv. NFS, and 0-4.5 equiv. NEt3 in the listed solvent with
1-1.5 atm each C3H6 and CO2 for 2 hrs.-3 days at room temperature. bNo gas added. c6 hrs. R.T., 1 hr. at
50 °C. MSA = methylsuccinic anhydride. Yield determined by GCMS using naphthalene as the standard.
184
Entry Ni Species [CO] [CO]
equiv.
[B] [B]
equiv.
Solvent MSA
(%)
1 B PF 3 NEt3 4 30:70
C6D6:
THF-d8
0.01
2 B PF 3 NEt3 4 50:50
C6D6:
THF-d8
0.03
3 B PF 3 NEt3 4 C6D6 0.05
4 B PF 9 NEt3 12 50:50
C6D6:
THF-d8
0.68
5 B PF 9 NEt3 24 50:50
C6D6:
THF-d8
1.60
6 B PF 9 NEt3 24 C6H6 2.00
Entry Ni
Species
[CO] [CO]
equiv.
[B] [B]
equiv.
Solvent Gas
Pressure
MSA
(%)
1 B PF 3 NEt3 4 50:50
C6D6:
THF-d8
1.5 atm
each
0.03
2 B PF 3 NEt3 4 50:50
C6D6:
THF-d8
5 atm
each
0.02
3 B PF 9 NEt3 12 50:50
C6D6:
THF-d8
1.5 atm
each
0.68
4 B PF 9 NEt3 12 50:50
C6D6:
THF-d8
None 0.02
Table A1.9: Effect of Gas in Headspacea
aReaction conditions: 5-21 mg Ni species, X equiv. PF, and Y equiv. NEt3 in the listed solvent with C3H6
and CO2 at 50-70 °C for up to 2 days. MSA = methylsuccinic anhydride. Yield determined by GCMS
using naphthalene as the standard.
Table A1.8: Solvent Screena
aReaction conditions: 6-21 mg Ni species, X equiv. [CO], and Y equiv. [B] in the listed solvent with 0.5-
1.5 atm each C3H6 and CO2 at 50-70 °C for up to 3 days. MSA = methylsuccinic anhydride. Yield
determined by GCMS using naphthalene as the standard.
Table A1.9: Effect of Gas in Headspacea
185
Entry Ni Species [CO] [CO]
equiv.
[B] [B]
equiv.
MSA
(%)
1 B TCPF 4 NEt3 4 4.80
2 B TCPF 4 NEt3 6 4.80
3 B TCPF 4 NEt3 8 5.20
4 B TCPF 4 NEt3 12 3.80
5 B TCPF 4 NEt3 16 7.30
6 B TCPF 4 NEt3 20 7.66
7 B TCPF 4 NEt3 40 7.47
8 B TCPF 4 NEt3 60 7.30
9 B TCPF 4 NEt3 100 4.60
10 B TCPF 4 NEt3 200 5.70
Entry Ni
Species
[CO] [CO]
equiv.
[B] [B]
equiv.
Gas Pressure MSA
(%)
1 B PF 3 NEt3 4 1.5 atm each 0.05
2 B PF 9 NEt3 12 None 0.09
3 B PF 9 NEt3 24 0.5 atm each 2.00
4 B TCPF 3 NEt3 4 1 atm each 5.60
5 B TCPF 4 NEt3 20 None 9.00
6 B TCPF 9 NEt3 24 0.5 atm each 5.40
7 B TCPF 9 NEt3 45 None 6.90
8 B TCPF 9 NEt3 100 0.5 atm each 4.00
9 B NFS 3 NEt3 4 0.5 atm each 6.30
10 B NFS 4 NEt3 10 1 atm each 9.00
11 B NFS 9 NEt3 24 0.5 atm each 7.40
Table A1.11: CO Surrogate Screena
aReaction conditions: 10 mg Ni species, X equiv. TCPF, and Y equiv. NEt3 in C6H6, no
gases, at room temperature for 2.5 hrs. MSA = methylsuccinic anhydride. Yield determined
by GCMS using naphthalene as the standard.
aReaction conditions: 8-10 mg Ni species, X equiv. [CO], and Y equiv. NEt3 in C6D6 or C6H6 with
indicated pressures of C3H6 and CO2 at 50-70 °C for up to 3 days. MSA = methylsuccinic anhydride.
Yield determined by GCMS using naphthalene as the standard.
Table A1.10: CO Surrogate:Base Ratio Screena
186
Entry Ni
Species
[CO] [CO]
equiv.
[B] [B]
equiv.
Gas Pressure MSA
(%)
1 B PF 9 NEt3 24 0.5 atm each 2.00
2 B TCFP 4 NEt3 20 None 9.00
3 B NFS 4 NEt3 10 1 atm each 9.00
4b B PF 9 DBU 12 1 atm each -
5 B TCPF 3 DBU 4 1 atm each 0.57
6 B NFS 3 DBU 4 1 atm each 0.77
7 B PF 9 DMAP 12 0.5 atm each -
8 B TCPF 9 DMAP 12 0.5 atm each 4.00
9b B NFS 9 DMAP 12 0.5 atm each 6.30
10 B TCPF 4 - - 1 atm each 5.20
Table A1.12: Base Screena
aReaction conditions: 8-10 mg Ni species, X equiv. of [CO], and Y equiv. [B] in C6D6 or C6H6 with
indicated pressures of C3H6 and CO2 at 50-70 °C for up to 3 days. bRoom temperature for 1-2 hrs. MSA
= methylsuccinic anhydride. Yield determined by GCMS using naphthalene as the standard.
187
Select NMR Spectra:
Figure A1.23: 31P{1H} NMR spectra showing attempts at reaction rate control using C6D6/THF-d8
solvent mixtures. 6-10 mg dcpe-Ni-lactone’, 3 equiv. PF, 4 equiv. NEt3. Blue = 12 hrs. at 70 °C, red = 6
hrs. at 50 °C, green = 24 hrs. at 50 °C, purple = 12 hrs. at 60 °C. B = dcpe-Ni-lactone’, * = dcpe-Ni(CO)2.
300 MHz spectrometer.
B B
*
*
*
*
B B
B
B B
B
C6D6
50:50 C6D6:THF-d8
50:50 C6D6:THF-d8
30:70 C6D6:THF-d8
188
Figure A1.24: 1H NMR spectra in C6D6 showing lack of decomposition of MSA under CO surrogate
reaction conditions over time. 10 mg MSA, 12 mg TCPF, 20 µl NEt3, 1 atm each C3H6 and CO2. * =
MSA peaks. # = NEt3 peaks. R.T. = room temperature. 600 MHz spectrometer.
* * * * #
30 min., R.T.
1.5 hrs., R.T.
1 hr., 60 °C
2 hrs., 60 °C
12 hrs., 60 °C
189
Figure A1.25: 1H NMR spectra in C6D6 showing lack of reaction of MSA with side products of CO
surrogate reactions. 10 mg MSA, 8 mg (1 equiv.) phenol, overnight at room temperature. * = MSA peaks.
# = phenol peaks. 500 MHz spectrometer. No evidence of reaction by GCMS.
*
* *
*
#
#
#
#
190
Figure A1.26: 1H NMR spectrum in C6D6 showing lack of reaction of MSA with side products of CO
surrogate reactions. 10 mg MSA, 8 mg (1 equiv.) phenol, 24 µl (2 equiv.) NEt3, overnight at room
temperature. * = MSA peaks. # = phenol peaks. o = NEt3 peaks. 500 MHz spectrometer. No evidence of
reaction by GCMS.
# #
# * *
*
o o
191
Preparation of dcpe-Ni(CO2):
1. [(Cy3P)2Ni]2(N2)
Make a suspension of 0.160 g Ni(acac)2 (acac = acetylacetonate) in ~5 mL toluene and
transfer to a 100-mL Schlenk flask. Dissolve 0.349 g PCy3 (2 equiv.) in ~5 mL toluene and
add to the Ni solution. Stir briefly, then cool to -35 °C for ~30 minutes. Inject 0.4 mL
Al(CH3)3 (1.2 equiv.) while cold and under a flow of N2, which will cause an immediate
color change to a deep reddish brown. While still under a positive N2 flow, allow the
reaction solution to warm slowly to room temperature with stirring. After ~2 hours, remove
the N2 flow and stir overnight. Remove the solvent under vacuum and recrystallize the
resulting reddish-brown oil from diethyl ether at -35 °C. Filter the resulting mixture and
wash the dark red crystalline product with cold diethyl ether. 31P{1H} NMR (C6D6, 300
MHz): 46.2 (s, 4P).
2. (Cy3P)2Ni(CO2)
Dissolve 0.095 g of the Ni-N2 dimer from (1) in ~2 mL toluene and transfer to a 25-mL
heavy-walled glass reaction vessel, giving a deep purple solution. Add ~1 atm CO2 at -196
°C using a calibrated gas bulb, resulting in an immediate solution color change to reddish
orange. Stir the mixture at room temperature for ~8 hours, which allows an orange
precipitate to accumulate. Filter the resulting mixture to yield an orange solid. 31P{1H}
NMR (C6D6, 300 MHz): 37.0 (s, 2P).
3. dcpe-Ni(CO2)
192
Add 0.010 g (Cy3P)2Ni(CO2) and 0.06 g dcpe (dry) to a J. Young NMR tube, then vacuum
transfer in ~0.5 mL C6D6 at -196 °C. Add ~2 atm CO2 using a calibrated gas bulb while
the mixture is still frozen. Thaw and mix. *Note: the reaction is immediate and goes to
completion. The product is not stable in solution and over the course of ~12 hours it
partially reverts back to the Ni-N2 dimer and partially decomposes. 31P{1H} NMR (C6D6,
300 MHz): 43.9 (s, 2P).
Figure A1.27: 31P{1H} NMR spectra in C6D6 showing the products of the stepwise synthesis of dcpe-
Ni(CO2). 300 MHz spectrometer. * = free PCy3.
[(Cy3P)2Ni]2(N2)
(Cy3P)2Ni(CO2)
dcpe-Ni(CO2)
*
193
1.2.B.: Nickel-Promoted Coupling of Isocyanates and Alkenes for the Production of
N-Substituted Acrylamides
General Considerations
All manipulations were carried out under a nitrogen or argon atmosphere using standard
Schlenk, vacuum, cannula, or glovebox techniques. The dcpe-Ni(COD) complex was
prepared according to previously published procedures.5,6 Bulk solvents were dried and
deoxygenated using literature procedures.3 All chemicals were purchased from VWR,
Aldrich, Fisher Scientific, Strem, or Cambridge Isotope Laboratories. NMR solvents were
dried over 3 Å molecular sieves and then used without further manipulation, or sodium and
then vacuum transferred prior to use. 1H, 13C, and 31P NMR spectra were recorded on
Bruker 300 MHz Avance II+, 300 MHz DRX, 500 MHz DRX or 600 MHz spectrometers
at ambient temperature, unless otherwise noted. Chemical shifts are reported in ppm; J
values are given in Hz. In some cases, the inequivalent 31P nuclei have coupling constants
too small to be observed due to the combination of 2J and 3J couplings and are therefore
denoted as singlets. 1H and 13C chemical shifts are referenced to residual solvent signals;
31P NMR chemical shifts are referenced to an external standard of H3PO4. Probe
temperatures were calibrated using ethylene glycol and methanol as previously described.4
GCMS analyses were performed using an Agilent gas chromatogram with mass
spectrometer. Product identities were confirmed by comparison to authentic samples or
previously reported spectra.9-14 All high-pressure coupling reactions were performed using
a Parr 5500 series compact reactor with glass insert.
*For procedure for dcpe-Ni(COD) preparation, see Appendix 1, section 1.2.A.ii.
194
Nickel-Promoted Coupling Reactions between Ethylene and Isocyanates
General NMR Procedure:
Dissolve 10-15 mg dcpe-Ni(COD) and 0-2 equivalents Lewis acid or base additive in C6D6
or THF-d8 and transfer to a J. Young (total reaction pressure <4 atm) or high-pressure (total
reaction pressure <13 atm) NMR tube. Add 1-2 atm ethylene, then 1-10 equivalents
isocyanate at -196 °C via calibrated gas bulb. Allow the reaction to warm to room
temperature, then heat anywhere from 22-70 °C for up to 2 days. Add ~1 atm CO gas to
release any bound organic products and analyze using NMR spectroscopy and/or GCMS
using a naphthalene standard (add 10 µl of a 0.1M naphthalene solution in the reaction
solvent to the reaction mixture and dilute to a total volume of 0.5 mL).
Adjusted procedure for phenyl isocyanate:
Due to its boiling point, this substrate could not undergo vacuum transfer. Instead, after
preparing the dcpe-Ni(COD) NMR solution, add 1 atm ethylene at -196 °C via calibrated
gas bulb, warm to room temperature, and bring into the glovebox. Freeze the solution in
the coldwell, inject 1-10 equivalents phenyl isocyanate using a microsyringe, and quickly
remove the NMR tube from the glovebox while keeping the solution frozen. Re-add 1-2
Scheme A1.6: Nickel-promoted ethylene/isocyanate coupling pathway.
195
atm ethylene, then warm the solution to room temperature. Proceed as indicated for other
isocyanates.
General Parr Procedure:
Dissolve 10-15 mg dcpe-Ni(COD) and 0-20 equivalents base additive in C6H6 or THF and
transfer to a Parr base. In the case of phenyl isocyanate, add 2-10 equivalents using a
microsyringe and seal the Parr reactor. For ethyl and tert-butyl isocyanate, vacuum transfer
2-10 equivalents into an ampule at -196 °C and then flame-seal the ampule. Carefully add
the sealed ampule to the dcpe-Ni(COD) solution in the Parr base in the glovebox, and then
seal the reactor. Outside the glovebox, pressurize the reactor to 300-1000 psi ethylene and
stir at 60 °C for 16 hours. Cool the reactor in an ice bath for ~45 minutes, then partially
vent the ethylene pressure to ~100 psi. Bring the Parr reactor back into the glovebox and
filter the product solution through Celite. Analyze by NMR spectroscopy and/or GCMS.
196
Ni Species EtN=C=O tBuN=C=O PhN=C=O
A 60.0 ppm (s) 60.0 ppm (s) 60.0 ppm (s)
B 63.3 ppm (s) 63.3 ppm (s) 63.3 ppm (s)
C 53.1 ppm
(d, J = 10)
74.6 ppm
(d, J = 10)
- -
D 48.5 ppm
(d, J = 8)
71.6 ppm
(d, J = 8)
- -
E 60.5 ppm
(d, J = 48)
70.9 ppm
(d, J = 48)
59.9 ppm
(d, J = 55)
68.2 ppm
(d, J = 55)
51.9 ppm
(d, J = 48)
73.8 ppm
(d, J = 49)
F 58.4 ppm (s)
65.8 ppm (s)
62.4 ppm (s)
68.6 ppm (s)
57.2 ppm (s)
66.5 ppm (s)
*C6D6, 300 or 300+ MHz spectrometers; see labeling guide A-F on the next page
Table A1.13: 31P{1H} NMR Data for Nickel-Bound Species
197
Select NMR Spectra:
NMR Labeling Guide-
A = dcpe-Ni(COD), B = dcpe-Ni-ethylene, C/D = Ni complexes with only isocyanate(s),
1. R = Et
dcpe-Ni(COD) + C2H4, 90 min., R.T.
+ Et-N=C=O
1 hr., R.T.
12 hrs., R.T.
3 hrs., 50 °C
12 hrs., 60 °C
12 hrs., 70 °C
A B
C C
E E
F F
Figure A1.28: 31P{1H} NMR spectra in C6D6 for the stepwise addition of 2 atm ethylene followed by 1
equiv. ethyl isocyanate. 300 MHz spectrometer.
E = dcpe-Ni-η2-acrylamide F = dcpe-Ni-metallocycle
198
Figure A1.30: 31P{1H} NMR spectra in C6D6 for the stepwise addition of 2 atm ethylene followed by 10
equiv. ethyl isocyanate, at room temperature for 2 days. 300 MHz spectrometer.
dcpe-Ni(COD) + C2H4, 30 min.
+ EtN=C=O
+ 2 hrs.
+ 12 hrs.
+ 48 hrs.
A B
C C
E E
D D
Figure A1.29: 1H NMR spectrum in C6D6 after CO gas addition to the ethylene/ethyl isocyanate coupling
reaction from Figure A1.28. Red box indicates potential acrylamide product, green indicates potential
succinimide. 500 MHz spectrometer.
199
Figure A1.31: 1H NMR spectrum in C6D6 for the confirmed N-ethyl acrylamide produced from nickel-
promoted ethylene/ethyl isocyanate coupling. (A) is the full spectrum, (B) is the expanded olefinic region.
Peak for –NH is likely too broad to observe. 500 MHz spectrometer.
A.
B.
-CH=CH2
-CH2
-CH3
200
A.
-CH2CH3
Figure A1.32: 1H NMR spectrum in C6D6 for the confirmed isolated N-ethyl succinimide produced from
nickel-promoted ethylene/ethyl isocyanate/CO coupling. (A) is the full spectrum, (B) is the expanded
characteristic ethyl –CH2 peak. Peaks for ethyl –CH3 and ring –CH2 groups are in the overlapping 1-2.1
ppm region. 500 MHz spectrometer.
B.
201
Figure A1.33: 31P{1H} NMR spectra in C6D6 for the coupling reaction of ethylene and ethyl isocyanate
(blue) versus the reaction of dcpe-Ni(COD) with N-ethyl acrylamide (red). Slightly shifted due to
reaction mixture, but matches the nickel η2-acrylamide complex. 300 MHz spectrometer.
E E
E
E
Figure A1.34: 31P{1H} NMR spectra in C6D6 for the coupling reaction of ethylene and ethyl isocyanate
using C2H4 (blue) versus 13C2H4 (red). Only the boxed peaks showed splitting changes due to
incorporation of 13C2H4. 300 MHz spectrometer.
dcpe-Ni(COD) + C2H4 + EtN=C=O
dcpe-Ni(COD) + H2C=CHC(O)NHEt
dcpe-Ni(COD) + C2H4 + EtN=C=O
dcpe-Ni(COD) + 13C2H4 + EtN=C=O
B
A
202
A.
B.
J = 12 Hz J = 30 Hz J = 57 Hz
J = 21 Hz
J = 57 Hz
J = 2 Hz
203
Figure A1.35: A. 31P{1H} and B. 13C{1H} NMR spectra (expanded sections, 1-3) in C6D6 for the
coupling reaction of ethylene and ethyl isocyanate using 13C2H4. Matching coupling constant denote
corresponding species. 31P: 300 MHz spectrometer, 13C: 500 MHz spectrometer.
J = 22 Hz
J = 12 Hz
J = 31 Hz
J = 32 Hz
204
Figure A1.36: 1HNMR spectrum in C6D6 indicating the ethyl isocyanate cyclotrimerization product. 500
MHz spectrometer.
Figure A1.37: 31P{1H} NMR spectrum in C6D6 for the coupling reaction between 8 atm ethylene and 3
equiv. Et-N=C=O. 300 MHz spectrometer.
B
C C
D D
205
Entry Et-N=C=O
(Equiv.)
C2H4 (psi) Products
1 5 300 η2-acrylamide: yes
Isocyanurate: yes
2 2 500 η2-acrylamide: no
Isocyanurate: no
3 10 500 η2-acrylamide: yes
Isocyanurate: yes
4b 10 500 η2-acrylamide: yes
Isocyanurate: yes
5 5 600 η2-acrylamide: yes
Isocyanurate: yes
6 10 750 η2-acrylamide: no
Isocyanurate: no
7 50 750 η2-acrylamide:
trace
Isocyanurate: yes
8c 50 750 η2-acrylamide:
trace
Isocyanurate: yes
9 2 1000 η2-acrylamide: no
Isocyanurate: no
10 5 1000 η2-acrylamide: no
Isocyanurate: no
11 10 1000 η2-acrylamide: no
Isocyanurate: no
Table A1.14: Nickel-promoted reaction of ethylene and ethyl isocyanate under high
pressure conditions.a
aReaction conditions: 10 mg dcpe-Ni(COD) in ~5 mL C6H6 with X equiv. EtN=C=O and Y psi of C2H4
at 60 °C for 16 hrs. bIn THF solvent. cWith 1 equiv. sodium 3-fluorophenoxide base.
206
E E
A
B
*
Figure A1.38: 31P{1H} NMR spectra in C6D6 for the Parr coupling reaction between 300 psi (blue) or
100 psi (red) ethylene and 5 or 2 equiv. Et-N=C=O, respectively. * = unidentified side product. 300 MHz
spectrometer.
Figure A1.39: 1H NMR spectra in C6D6 for the Parr coupling reaction between 300 psi (blue) or 100 psi
(red) ethylene and 5 or 2 equiv. Et-N=C=O, respectively. Boxes indicate peaks for the isocyanurate. 500
MHz spectrometer.
207
2. R = tBu
Free
COD
Free
COD
Free tBuNCO
-CH=CH2
Figure A1.40: 1H NMR spectrum in C6D6 for the coupling reaction between 1 atm ethylene and 10 equiv. tBu-N=C=O, after CO addition. 500 MHz spectrometer.
E E F
C
B
dcpe-Ni(COD) + C2H4 + tBuN=C=O
dcpe-Ni(COD) + tBuN=C=O
Figure A1.41: 31P{1H} NMR spectra in C6D6 for the coupling reaction between 1 atm ethylene and 10
equiv. tBu-N=C=O with simultaneous (blue) or stepwise (red and green) reagent additions. 300 MHz
spectrometer.
208
3. R = Ph
* * *
Free
COD Free
COD
Free
PhNCO
Figure A1.42: 1H NMR spectrum in C6D6 for the coupling reaction between 1.5 atm ethylene and 2
equiv. Ph-N=C=O, after CO addition. Green box indicates potential succinimide, the red box shows lack
of vinylic peaks. 600 MHz spectrometer.
F F
dcpe-Ni(COD) + C2H4 + PhN=C=O
dcpe-Ni(COD) + C2H4 + PhN=C=O + NaBAr4F, immediate
dcpe-Ni(COD) + C2H4 + PhN=C=O + NaBAr4F, 1 hr. at 60 °C
Figure A1.43: 31P{1H} NMR spectra in C6D6 for the coupling reaction between 1.5 atm ethylene and 2
equiv. Ph-N=C=O, without (blue) and with (red and green) 1 equiv. Lewis acid. New species in red and
green spectra did not yield acrylamide. Suspected C/D-type complexes. 300 MHz spectrometer.
209
Figure A1.44: 1H NMR spectrum in C6D6 for the coupling reaction between 1 atm ethylene and 2 equiv.
Ph-N=C=O in the presence of 1 equiv. Lewis acid, after CO addition. No acrylamide generated. 500 MHz
spectrometer.
F F
dcpe-Ni(COD) + C2H4 + PhN=C=O
dcpe-Ni(COD) + C2H4 + PhN=C=O + NEt3
12 hrs., 60 °C
Figure A1.45: 31P{1H} NMR spectra in C6D6 for the coupling reaction between 1 atm ethylene and 2
equiv. Ph-N=C=O, without (blue) and with (red) 2 equiv. NEt3. 300 MHz spectrometer.
210
dcpe-Ni(COD) + C2H4 + PhN=C=O
+ DBU, 20 hrs., 60 °C
F F
B
Figure A1.46: 31P{1H} NMR spectra in THF-d8 for the coupling reaction between 1 atm ethylene and 2
equiv. Ph-N=C=O, before (blue) and after (red) addition of 2 equiv. DBU. New species in red did not
yield acrylamide. 300 MHz spectrometer.
Figure A1.47: 1H NMR spectrum in THF-d8 for the coupling reaction between 1 atm ethylene and 2
equiv. Ph-N=C=O in the presence of 2 equiv. DBU, after CO addition. Olefinic region. 2D NMR
indicates the boxed peaks are not acrylamide. 600 MHz spectrometer.
211
B
A
F F
dcpe-Ni(COD) + C2H4 + PhN=C=O
+ NaOFPh, 12 hrs., 60 °C
Figure A1.48: 31P{1H} NMR spectra in THF-d8 for the coupling reaction between 1 atm ethylene and 4
equiv. Ph-N=C=O, before (blue) and after (red) addition of 2 equiv. NaOFPh (sodium 3-
fluorophenoxide). New species in red did not yield acrylamide. 300 MHz spectrometer.
Figure A1.49: 1H NMR spectrum in THF-d8 for the coupling reaction between 1 atm ethylene and 4
equiv. Ph-N=C=O in the presence of 2 equiv. sodium 3-fluorophenoxide, after CO addition. 2D NMR
indicates the boxed peaks are not acrylamide. 600 MHz spectrometer.
212
F F
B
dcpe-Ni(COD) + C2H4 + PhN=C=O
+ NaOtBu, 12 hrs., 60 °C
Figure A1.50: 31P{1H} NMR spectra for the coupling reaction between 1 atm ethylene and 2 equiv. Ph-
N=C=O, before (blue) and after (red) addition of 2 equiv. NaOtBu. Blue = C6D6, red = THF-d8. 300 MHz
spectrometer.
A.
213
C.
B.
Figure A1.51: 1H NMR (A), HSQC (B), and COSY (C) spectra in THF-d8 for the coupling reaction
between 1 atm ethylene and 2 equiv. Ph-N=C=O in the presence of 2 equiv. NaOtBu. Red boxes indicated
acrylamide vinylic peaks. 500 MHz spectrometer.
-CH=CH2 -CH=CH2
214
1.3: Transition-Metal-Catalyzed Hydrogenation of CO2 to Methanol
General Considerations
All manipulations were carried out under a nitrogen or argon atmosphere using standard
Schlenk, vacuum, cannula, or glovebox techniques. Catalysts 1, 2, and 1-formate were
prepared as previously described.15,16 All other chemicals were purchased from Aldrich,
Fisher, Strem, Oakwood Chemicals, VWR, or Cambridge Isotope Laboratories. Liquid
amine substrates and 1,8-diazabicyclo[5.4.0]undec-7-ene (DBU) were dried over calcium
hydride or sodium hydride, purified by vacuum transfer or distillation, and stored over 3 Å
molecular sieves. Lithium triflate (LiOTf) was gently heated overnight under vacuum to
remove trace water. Molecular sieves (3 Å) for catalytic reactions were rigorously dried
under vacuum for 48-72 hours. Bulk solvents were dried and deoxygenated using literature
procedures.3 NMR solvents were dried over 3 Å molecular sieves and then used without
further manipulation, or sodium and then vacuum transferred prior to use. Hydrogen and
carbon dioxide were purchased from Airgas and were used as received. 1H, 13C and 31P
NMR spectra were recorded on Bruker 300 MHz Avance II+, 300 MHz DRX, 500 MHz
DRX or 600 MHz spectrometers at ambient temperature, unless otherwise noted. Chemical
shifts are reported in ppm; J values are given in Hz. 1H and 13C chemical shifts are
referenced to residual solvent signals; 31P chemical shifts are referenced to an external
standard of H3PO4. Probe temperatures were calibrated using ethylene glycol and methanol
as previously described.4 Gas chromatography was performed on a Thermofisher Scientific
Trace 1300 Series gas chromatograph with a FID or on an Agilent gas chromatogram with
mass spectrometer, both using helium as the carrier gas. All high-pressure CO2
hydrogenation reactions were performed using a Parr 5500 series compact reactor with
215
glass insert. X-ray crystallographic data was collected on a Bruker SMART CCD or a
Bruker Prospector diffractometer with Apex II detectors. Samples were collected in inert
oil and quickly transferred to a cold gas stream. The structures were solved from direct
methods and Fourier syntheses, then refined by full-matrix least-squares procedures with
anisotropic thermal parameters for all non-hydrogen atoms. Crystallographic calculations
were carried out using SHELXTL.
216
General Procedure for Catalytic Reactions:
Step A: Inside of a glovebox, weigh out 2.02 g of 3Å molecular sieves and 21 mmol
morpholine into a glass Parr insert. Add 8 mL solvent and 15 µmol catalyst (300 µl of a
0.5M solution), seal the reactor, and remove from the glovebox. Pressurize with the desired
amount of CO2, then H2. Heat at the optimized temperature with stirring for 16 hours. Cool
the reactor in an ice bath for 45 minutes, then partially vent to 100 psi. Bring the reactor
back inside of the glovebox, fully vent it, and open it. Filter the reaction solution, and then
analyze an aliquot by GC-FID using mesitylene as a standard (100 µl of reaction solution
with 0.024M or 0.0024M mesitylene standard after final dilution to 1 mL total volume).
Keep the remaining step A solution in the glovebox in preparation for step B.
Step B: Inside of a glovebox, weigh out 1 mmol LiOTf and 2.5 mmol DBU into a glass
Parr insert. Add the volume of reaction solution from step A corresponding to the desired
amount of 4-formylmorpholine for step B and dilute with solvent to a final volume of 5
mL. Inject 10 µmol catalyst (100 µl of a 0.5M solution) and seal the reactor. Outside of the
glovebox, pressurize the reactor with the same amount of H2 as in step A and heat to the
same temperature (with stirring) for 16 hours. Cool the reactor in an ice bath for 45 minutes,
then vent and open. Filter the reaction solution and then analyze by GC-FID using a
mesitylene standard (100 µl of reaction solution with 0.024M or 0.0024M mesitylene
standard after final dilution to 1 mL total volume).
217
Synthesis Procedure for Generating Morphylammonium Morphylcarbamate:
Transfer morpholine and tetrahydrofuran (THF) into a glass Parr insert and seal the reactor.
Pressurize with 1000 psi CO2 and stir for 16 hours at 120 °C. Cool the reactor in an ice
bath for 45 minutes, then open and quickly filter out the resulting white solid as the product.
X-ray quality crystals were obtained from a highly concentrated solution in 1,4-dioxane.
Figure A1.52: Molecular structure of morphylammonium morphylcarbamate with ellipsoids at 50%
probability. Select bond distances, Å: (C7)-(O1) = 1.27, (C7)-(O2) = 1.26, (C7)-(N1) = 1.39.
218
Reaction Optimization Tables For:
Entryb Substrate
(mmol)
CO2/H2
(psi)
Catalyst
(µmol)
FM
(mmol)
CH3OH
(mmol)
TONd Additive
1Ac Morph.
(21)
100/900 15 4.87 - 325 Mol.
sieves
1Bc FM
(1.25)
0/900 10 0.44 0.81 81 -
2A Morph.
(21)
100/900 15 12.57 - 838 -
2B FM
(1.25)
0/900 10 1.20 0.14 14 -
3A Morph.
(21)
100/900 15 12.49 - 832 -
3B FM
(1.25)
0/900 10 1.48 - - Mol.
sieves
4A Morph.
(21)
100/900 15 4.84 - 322 Mol.
sieves
4B FM
(1.25)
0/900 10 1.00 - - Mol.
sieves
Table A1.15: Effect of Molecular Sievesa
aReaction conditions: Step A: 21 mmol morpholine with or without 2.02 g molecular sieves, 15 µmol
catalyst 1 (0.07 mol%) in 8 mL dioxane, 100 psi CO2, 900 psi H2, 120 °C, 16 hrs. Step B: 1.25 mmol
FM from step A solution, 10 µmol catalyst 1 (0.8 mol%), 1 mmol LiOTf, 2.5 mmol DBU, with or
without 2.02 g molecular sieves in 5 mL total volume of solution, 100 psi H2, 120 °C, 16 hrs. bEach
entry represents only one trial unless otherwise indicated. cAverage of two trials. dTurnover number
determined by GC-FID analysis of formamide in step A and methanol in step B.
219
Entryb Substrate
(mmol)
CO2/H2
(psi)
Catalyst
(µmol)
FM
(mmol)
CH3OH
(mmol)
TONd LiOTf
(mmol)
1Ac Morph.
(21)
100/900 15 4.87 - 325 -
1Bc FM
(1.25)
0/900 10 0.44 0.81 81 1
5A Morph.
(21)
100/900 15 5.23 - 348 -
5B FM
(1.25)
0/900 10 1.17 0.08 8 -
6A Morph.
(21)
100/900 15 1.86 - 124 1
6B FM
(1.25)
0/900 10 0.28 1.07 107 1
Table A1.16: Effect of LiOTfa
aReaction conditions: Step A: 21 mmol morpholine, 2.02 g molecular sieves, 15 µmol catalyst 1 (0.07
mol%) in 8 mL dioxane, with or without 1 mmol LiOTf, 100 psi CO2, 900 psi H2, 120 °C, 16 hrs. Step
B: 1.25 mmol FM from step A solution, 10 µmol catalyst 1 (0.8 mol%), with or without 1 mmol LiOTf,
2.5 mmol DBU in 5 mL total volume of solution, 100 psi H2, 120 °C, 16 hrs. bEach entry represents
only one trial unless otherwise indicated. cAverage of two trials. dTurnover number determined by GC-
FID analysis of formamide in step A and methanol in step B.
220
Entryb Substrate
(mmol)
CO2/H2
(psi)
Catalyst
(µmol)
FM
(mmol)
CH3OH
(mmol)
TONd DBU
(mmol)
1Ac Morph.
(21)
100/900 15 4.87 - 325 -
1Bc FM
(1.25)
0/900 10 0.44 0.81 81 2.5
7A Morph.
(21)
100/900 15 5.02 - 334 -
7B FM
(1.25)
0/900 10 1.38 - - -
8Ac Morph.
(21)
100/900 15 11.90 - 793 2.5
8Bc FM
(1.25)
0/900 10 0.32 0.55 55 2.5
9Ac Morph.
(21)
100/900 15 4.80 - 320 -
9Bc FM
(1.25)
0/900 10 0.01 0.84 84 5.0
10Ac Morph.
(21)
100/900 15 9.16 - 610 5.0
10Bc FM
(1.25)
0/900 10 1.77 - - -
Table A1.17: Effect of DBUa
aReaction conditions: Step A: 21 mmol morpholine, 2.02 g molecular sieves, 15 µmol catalyst 1 (0.07
mol%) in 8 mL dioxane, with or without 2.5-5 mmol DBU, 100 psi CO2, 900 psi H2, 120 °C, 16 hrs.
Step B: 1.25 mmol FM from step A solution, 10 µmol catalyst 1 (0.8 mol%), 1 mmol LiOTf, with or
without 2.5-5 mmol DBU in 5 mL total volume of solution, 100 psi H2, 120 °C, 16 hrs. bEach entry
represents only one trial unless otherwise indicated. cAverage of two trials. dTurnover number
determined by GC-FID analysis of formamide in step A and methanol in step B.
221
Entryb Substrate
(mmol)
CO2/H2
(psi)
Catalyst
(µmol)
FM
(mmol)
CH3OH
(mmol)
TONd Carbamate
(mmol)
1Ac Morph.
(21)
100/900 15 4.87 - 325 -
1Bc FM
(1.25)
0/900 10 0.44 0.81 81 2.5
11A Morph.
(21)
100/900 15 5.94 - 396 -
11B FM
(1.25)
0/900 10 1.21 0.04 4 1.25
12Ae Carbamate
(3)
100/900 15 0.64 - 42
12Be Carbamate
(1.25)
0/900 10 0.11 - -
Table A1.18: Effect of Morphylammonium Morphylcarbamatea
aReaction conditions: Step A: 21 mmol morpholine, 2.02 g molecular sieves, 15 µmol catalyst 1 (0.07
mol%) in 8 mL dioxane, 100 psi CO2, 900 psi H2, 120 °C, 16 hrs. Step B: 1.25 mmol FM from step A
solution, 10 µmol catalyst 1 (0.8 mol%), 1 mmol LiOTf, 2.5 mmol DBU, with or without 1.25 mmol
morphyl carbamate in 5 mL total volume of solution, 100 psi H2, 120 °C, 16 hrs. bEach entry represents
only one trial unless otherwise indicated. cAverage of two trials. dTurnover number determined by GC-
FID analysis of formamide in step A and methanol in step B. eCarbamate substrate in both steps.
222
Entryb Substrate
(mmol)
CO2/H2
(psi)
Catalyst
(µmol)
FM
(mmol)
CH3OH
(mmol)
TONc Solvent
1A Morph.
(21)
100/900 15 4.42 - 294 THF
1B FM
(1.25)
0/900 10 0.01 1.13 113
2A Morph.
(21)
100/900 15 4.30 - 286 Dioxane
2B FM
(1.25)
0/900 10 0.03 1.14 113
3A Morph.
(21)
100/900 15 3.30 - 220 Toluene
3B FM
(1.25)
0/900 10 0.01 1.29 129
4A Morph.
(21)
100/900 15 3.37 - 224 Ethyl
Acetate
4B FM
(1.25)
0/900 10 0.09 0.37 37
Table A1.19: Solvent Screena
aReaction conditions: Step A: 21 mmol morpholine, 2.02 g molecular sieves, 15 µmol catalyst 1 (0.07
mol%) in 8 mL solvent, 100 psi CO2, 900 psi H2, 120 °C, 16 hrs. Step B: 1.25 mmol FM from step A
solution, 10 µmol catalyst 1 (0.8 mol%), 1 mmol LiOTf, 2.5 mmol DBU in 5 mL total volume of
solution, 100 psi H2, 120 °C, 16 hrs. bAverage of two trials unless otherwise indicated. cTurnover
number determined by GC-FID analysis of formamide in step A and methanol in step B.
223
Entryb Substrate
(mmol)
CO2/H2
(psi)
Catalyst
(µmol)
FM
(mmol)
CH3OH
(mmol)
TONd Max.
TON
1A Morph.
(21)
100/900 15 4.42 - 294 1400
1B FM
(1.25)
0/900 10 0.01 1.13 113 125
5A Morph.
(21)
200/800 15 6.91 - 460 1400
5B FM
(1.25)
0/800 10 0.01 1.08 108 125
6Ac Morph.
(21)
200/800 15 6.86 - 457 1400
6Bc FM
(3.05)
0/800 10 0.25 2.34 234 305
7Ac Morph.
(21)
200/1200 15 9.44 - 629 1400
7Bc FM
(4.20)
0/1200 10 0.10 3.52 352 420
8A Morph.
(21)
200/1200 15 9.24 - 616 1400
8B FM
(4.00)
0/1200 10 0.05 3.40 340 400
Table A1.20: Pressure Screena
aReaction conditions: Step A: 21 mmol morpholine, 2.02 g molecular sieves, 15 µmol catalyst 1 (0.07
mol%) in 8 mL THF, X psi CO2, Y psi H2, 120 °C, 16 hrs. Step B: 1.25-4 mmol FM from step A
solution, 10 µmol catalyst 1 (0.25-0.8 mol%), 1 mmol LiOTf, 2.5 mmol DBU in 5 mL total volume of
solution, Y psi H2, 120 °C, 16 hrs. bAverage of two trials unless otherwise indicated. cOnly one trial. dTurnover number determined by GC-FID analysis of formamide in step A and methanol in step B.
224
Entryb Substrate
(mmol)
CO2/H2
(psi)
Catalyst
(µmol)
FM
(mmol)
CH3OH
(mmol)
TONc LiOTf,DBU
(mmol)
8A Morph.
(21)
200/1200 15 9.24 - 616 -
8B FM
(4.00)
0/1200 10 0.05 3.40 340 1, 2.5
9A Morph.
(21)
200/1200 15 9.19 - 612 -
9B FM
(4.00)
0/1200 10 0.20 3.27 327 1, 4.0
10A Morph.
(21)
200/1200 15 9.29 - 619 -
10B FM
(4.00)
0/1200 10 0.09 3.41 341 1.5, 2.5
11A Morph.
(21)
200/1200 15 9.28 - 618 -
11B FM
(4.00)
0/1200 10 0.12 3.20 320 1.5, 4.0
Table A1.21: LiOTf/DBU Screena
aReaction conditions: Step A: 21 mmol morpholine, 2.02 g molecular sieves, 15 µmol catalyst 1 (0.07
mol%) in 8 mL THF, 200 psi CO2, 1200 psi H2, 120 °C, 16 hrs. Step B: 4 mmol FM from step A
solution, 10 µmol catalyst 1 (0.25 mol%), 1-1.5 mmol LiOTf, 2.5-4 mmol DBU in 5 mL total volume
of solution, 1200 psi H2, 120 °C, 16 hrs. bAverage of two trials unless otherwise indicated. cTurnover
number determined by GC-FID analysis of formamide in step A and methanol in step B.
225
Entryb Substrate
(mmol)
CO2/H2
(psi)
Catalyst
(µmol)
FM
(mmol)
CH3OH
(mmol)
TONd Catalyst
8A Morph.
(21)
200/1200 15 9.24 - 616 1
8B FM
(4.00)
0/1200 10 0.05 3.40 340 1
12A Morph.
(21)
200/1200 15 3.20 - 213 2
12B FM
(4.00)
0/1200 10 - - - NA
13Ac Morph.
(21)
500/500 15 2.94 - 196 2
13Bc FM
(4.00)
0/500 10 0.01 1.29 129 NA
14Ac Morph.
(21)
500/500 5 1.67 - 334 2
14Bc FM
(4.00)
0/500 10 - - - NA
Table A1.22: Catalyst Screena
aReaction conditions: Step A: 21 mmol morpholine, 2.02 g molecular sieves, 5-15 µmol catalyst 1 or 2
in 8 mL THF, 200 psi CO2, 1200 psi H2 (or 500:500), 120 °C, 16 hrs. Step B: 4 mmol FM from step A
solution, 10 µmol catalyst 1 (0.25 mol%), 1 mmol LiOTf, 2.5 mmol DBU in 5 mL total volume of
solution, 1200 psi H2, 120 °C, 16 hrs. bAverage of two trials unless otherwise indicated. cOnly one trial. dTurnover number determined by GC-FID analysis of formamide in step A and methanol in step B.
226
Entryb Substrate
(mmol)
CO2/H2
(psi)
Catalyst
(µmol)
FM
(mmol)
CH3OH
(mmol)
TONc Temp.
(°C)
8A Morph.
(21)
200/1200 15 9.24 - 616 120
8B FM
(4.00)
0/1200 10 0.05 3.40 340 120
15A Morph.
(21)
200/1200 15 16.76 - 1117 100
15B FM
(4.00)
0/1200 10 0.01 3.60 360 100
16A Morph.
(21)
200/1200 15 15.24 - 1016 80
16B FM
(4.00)
0/1200 10 0.44 3.22 322 80
Table A1.23: Temperature Screena
aReaction conditions: Step A: 21 mmol morpholine, 2.02 g molecular sieves, 15 µmol catalyst 1 (0.07
mol%) in 8 mL THF, 200 psi CO2, 1200 psi H2, Z °C, 16 hrs. Step B: 4 mmol FM from step A solution,
10 µmol catalyst 1 (0.25 mol%), 1 mmol LiOTf, 2.5 mmol DBU in 5 mL total volume of solution,
1200 psi H2, Z °C, 16 hrs. bAverage of two trials unless otherwise indicated. cTurnover number
determined by GC-FID analysis of formamide in step A and methanol in step B.
227
Appendix 1 References
1 Crépy, K.V.L.; Imamoto, T. “PREPARATION OF (S,S)-1,2-BIS(tert-BUTYLMETHYLPHOSPHINO)
ETHANE ((S,S)-t-Bu-BISP*) AS A RHODIUM COMPLEX,” Organic Syntheses, 2005, 82, 22-29.
2 Yakelis, N.A.; Bergman, R.G. “Safe Preparation and Purification of Sodium Tetrakis[(3,5-
trifluoromethyl)phenyl]borate (NaBArF24): Reliable and Sensitive Analysis of Water in Solutions of
Fluorinated Tetraarylborates,” Organometallics, 2005, 24, 3579-3581.
3 Pangborn, A.B.; Giardello, M.A.; Grubbs, R.H.; Rosen, R.K.; Timmers, F.J. “Safe and Convenient
Procedure for Solvent Purification,” Organometallics, 1996, 15, 1518-1520.
4 Sandström, J. Dynamic NMR Spectroscopy; Academic Press: New York, 1982.
5 Hoberg, H.; Schaefer, D. “NICKEL(0)-INDUZIERTE C-C-VERKNÜPFUNG ZWISCHEN
KOHLENDIOXID UND ETHYLEN SOWIE MONO-ODER DI-SUBSTITUIERTEN ALKENEN,” J.
Organomet. Chem., 1983, 251, C51-C53.
6 Jin, D.; Williard, P.G.; Hazari, N.; Bernskoetter, W.H. “Effect of Sodium Cation on Metallocycle β-Hydride
Elimination in CO2-Ethylene Coupling to Acrylates,” Chem. Eur. J., 2014, 20, 3205-3211.
7 Jolly, P.W.; Jonas, K.; Krüger, C.; Tsay, Y.-H. “THE PREPARATION, REACTIONS, AND STRUCTURE
OF BIS[BIS-(TRICYCLOHEXYLPHOSPHINE)NICKEL] DINITROGEN, {[(C6H11)3P]2Ni}2N2,” J.
Organomet. Chem., 1971, 33, 109-122.
8 Aresta, M.; Nobile, C.F.; Albano, V.G.; Forni, E.; Manassero, M. “New Nickel-Carbon Dioxide Complex:
Synthesis, Properties, and Crystallographic Characterization of (Carbon dioxide)-
bis(tricyclohexylphosphine)nickel,” J.C.S. Chem. Comm., 1975, 636-637.
9 Maier, L. “The Direct Synthesis of Tris(N-Substituted Carbamoylethyl) Phosphine Oxides,” Helvetica
Chimica Acta, 1973, 56, 1252-1257.
10 Krivec, M.; Gazvoda, M.; Kranjc, K.; Polanc, S.; Kočevar, M. “A Way to Avoid Using Precious Metals:
The Application of High-Surface Activated Carbon for the Synthesis of Isoindoles via the Diels-Alder
Reaction of 2H-Pyran-2-ones,” J. Organic Chem., 2012, 77, 2857-2864.
11 Dokli, I.; Gredičak, M.; “Mechanochemical Ritter Reaction: A Rapid Approach to Functionalized Amides
at Room Temperature,” Chem. Eur. J., 2015, 2727-2732.
12 Garad, D.N.; Tanpure, S.D.; Mhaske, S.B. “Radical-mediated dehydrative preparation of cyclic imides
using (NH4)2S2O8-DMSO: application to the synthesis of vernakalant,” Beilstein J. Org. Chem., 2015, 11,
1008-1016.
13 Tang, X.-J.; Dolbier, W.R. “Efficient Cu-Catalyzed Atom Transfer Radical Addition Reactions of
Fluoroalkylsulfonyl Chlorides with Electron-deficient Alkenes Induced by Visible Light,” Angew. Chem. Int.
Ed., 2015, 54, 4246-4249.
14 Mo, F.; Yan, J.M.; Qiu, D.; Li, F.; Zhang, Y.; Wang, J. “Gold-Catalyzed Halogenation of Aromatics by N-
Halosuccinimides,” Angew. Chem. Int. Ed., 2010, 49, 2028-2032.
228
15 Bielinski, E.A.; Lagaditis, P.O.; Zhang, Y.; Mercado, B.Q.; Würtele, C.; Bernskoetter, W.H.; Hazari, N.;
Schneider, S. “Lewis Acid-Assisted Formic Acid Dehydrogenation Using a Pincer-Supported Iron Catalyst,”
J. Am. Chem. Soc., 2014, 136, 10234-10237.
16 Werkmeister, S.; Junge, K.; Wendt, B.; Alberico, E.; Jiao, H.; Baumann, W.; Junge, H.; Gallou,F.; Beller,
M. “Hydrogenation of Esters to Alcohols with a Well-Defined Iron Complex,” Angew. Chem. Int. Ed., 2014,
53, 8722-8726.
229
Appendix 2
2.2: Iron-Catalyzed Amide Formation from the Dehydrogenative Coupling of
Alcohols and Secondary Amines
General Considerations
All manipulations were carried out under a nitrogen or argon atmosphere using standard
Schlenk, vacuum, cannula, or glovebox techniques. Catalysts 1 and
(iPrPNHP)Fe(H)(HCO2)(CO) (1-formate) were prepared as previously described.1 Amide
products (for gas chromatography response factors) that were not commercially available
were prepared using previously reported procedures.2,3 All other chemicals were purchased
from Aldrich, Fisher, Strem, Synthonix, Oakwood Chemicals, VWR, or Cambridge
Isotope Laboratories. Liquid amine and alcohol substrates were dried over calcium hydride
or sodium hydride, purified by vacuum transfer or distillation, and stored over 3 Å
molecular sieves. Solid substrates were purified by sublimation, followed by
recrystallization (if necessary). Bulk solvents were dried and deoxygenated using literature
procedures.4 NMR solvents were dried over 3 Å molecular sieves and then used without
further manipulation, or sodium and then vacuum transferred prior to use. Hydrogen and
carbon dioxide were purchased from Airgas and were used as received. 1H, 13C and 31P
NMR spectra were recorded on Bruker 300 MHz Avance II+, 300 MHz DRX, 500 MHz
DRX or 600 MHz spectrometers at ambient temperature, unless otherwise noted. Chemical
shifts are reported in ppm; J values are given in Hz. 1H and 13C chemical shifts are
referenced to residual solvent signals; 31P chemical shifts are referenced to an external
standard of H3PO4. Probe temperatures were calibrated using ethylene glycol and methanol
230
as previously described.5 Gas chromatography was performed on a Thermofisher Scientific
Trace 1300 Series gas chromatograph with a FID using helium as the carrier gas.
General Procedure for the Formation of Complex 2
In a dry box, 5 mg of complex 1 was dissolved in either C6D6 or THF-d8 in a 20-mL
scintillation vial and then transferred to a J-Young tube. The tube was then sealed, removed
from the glovebox, degassed, and 6 equivalents of methanol or 13C-labeled methanol were
added via calibrated gas bulb.
Select NMR data for 2 (at -80 °C) 1H NMR (THF-d8, 300 MHz): -24.3 (t, 1H, J = 52, Fe-
H); 13C{1H} NMR (THF-d8, 300 MHz): 59.58 (s, Fe-O-CH3); 13C NMR (THF-d8, 300
MHz): 59.57 (q, J = 132, Fe-O-CH3);31P{1H} NMR (THF-d8, 300 MHz): 93.21 (d, J = 32).
General Methods for Catalytic Alcohol Dehydrogenation in the Presence of Amines
In a dry box, a 100-mL Schlenk tube was loaded with 5 mL of tetrahydrofuran (THF), 12
mmol of amine, 3 µmol of catalyst, and 3 mmol of alcohol, then sealed. It was immediately
placed in an oil bath preheated to 80 °C and stirred for 8 hours. It was then cooled in an ice
bath for 30 minutes prior to analysis. If analyzed using NMR spectroscopy, 100 µl of
reaction solution were added to an NMR tube with 395 µl CDCl3 and 5 µl of mesitylene
standard and an NMR delay time of 60 seconds was used. If analyzed by GC, 100 µl of
reaction solution was diluted to 1 mL with THF and a mesitylene standard (0.024 M or
0.0024M after final dilution) was added.
231
General Procedure for H2 Collection Studies
In a dry box, a 100-mL Schlenk flask was loaded with 2.5 mL THF, 6 mmol of morpholine,
1.5 µmol of catalyst 1, and 1.5 mmol methanol. The flask was fitted with a reflux condenser
and an adaptor with a stopcock, removed from the glovebox, and connected to a gas burette
setup that had been pre-sparged with N2. The connecting hoses/trap were subjected to two
evacuation/refill cycles, the trap was cooled in a dry ice/acetone bath, and a small amount
of vacuum was used to reset the starting water volume. The Schlenk flask was lowered into
an oil bath (preheated to 80 °C) and the system was allowed to equilibrate for 3.5 minutes,
after which a small amount of vacuum was used to restore the water volume in the buret
(with the connection to the reflux setup closed). The stopcock was then opened, the reaction
was allowed to proceed for 8 hours, and the change in the water level in the gas burette
was used to determine the TON. The procedure for determining TON has been previously
reported.1
232
Reaction Optimization Tables for:
Entry Solvent TON
1 Acetonitrile 30
2 Ethyl Acetate 49b
3 Toluene 297b
4 Dioxane 344b
5 Tetrahydrofuran, 10 mL 392
6 Tetrahydrofuran, 5 mL 503
Table A2.1: Reaction Solventa
aReaction conditions: 3 µmol catalyst 1 (0.1 mol%), 3 mmol
alcohol, and 12 mmol amine in 10 mL solvent at 80 °C for 8
hrs. Each entry is an average of two trials unless otherwise
indicated. TON determined by NMR analysis of the production
of amide. bOnly one trial.
233
Entry Time (hrs) TON
1 1 176b
2 4 383
3 8 503
4 16 553
Entry Temp. (°C) TON
1 R.T. 10
2 60 172
3 80 503
4 100 328
5 120 341
aReaction conditions: 3 µmol catalyst 1 (0.1 mol%), 3
mmol alcohol, and 12 mmol amine in 5 mL THF at Y °C
for 8 hrs. Each entry is an average of two trials unless
otherwise indicated. TON determined by NMR analysis
of the production of amide.
Table A2.3: Reaction Temperaturea
Table A2.2: Reaction Timea
aReaction conditions: 3 µmol catalyst 1 (0.1
mol%), 3 mmol alcohol, and 12 mmol amine in
5 mL THF at 80 °C for Z hrs. Each entry is an
average of two trials unless otherwise indicated.
TON determined by NMR analysis of the
production of amide. bAverage of three trials.
234
Entry CH3OH
(mmol)
Morpholine
(mmol)
CH3OH:
Morpholine
TON (FM)
1 3 3 1:1 261
2 3 12 1:4 503
3 3 30 1:10 425
4 12 3 4:1 163
Entry Catalyst TON
1 None 0b
2 FeCl2 0b
3 1-formate 0c
4 1 503
Table A2.5: Catalyst Control Experimentsa
aReaction conditions: 3 µmol catalyst 1 (0.1 mol%), A mmol alcohol, and B mmol amine in 5 mL
THF at 80 °C for 8 hrs. Each entry is an average of two trials unless otherwise indicated. TON
determined by NMR analysis of the production of amide.
aReaction conditions: 3 µmol catalyst X (0.1 mol%),
3 mmol alcohol, and 12 mmol amine in 5 mL THF at
80 °C for 8 hrs. Each entry is an average of two trials
unless otherwise indicated. TON determined by NMR
analysis of the production of amide. bOnly one trial.
cWith 10 mol% LiBF4 or LiOTf.
Table A2.4: Alcohol/Amine Ratioa
235
Entry Amine Additive (mol%) TON
1 Morpholine LiOTf, 10 370b
2 Morpholine None 503
3 Morpholine KOtBu, 1 216b
4 Morpholine KOtBu, 15 86b
5 Diethylamine None 176
6 Diethylamine KOtBu, 1 11
Table A2.6: Effect of Base or Lewis Acid Additivea
aReaction conditions: 3 µmol catalyst 1 (0.1 mol%), 3 mmol alcohol, 12 mmol amine, and X
mol% base (potassium tert-butoxide) or Lewis acid (lithium trifluoromethanesulfonate) in 5
mL THF at 80 °C for 8 hrs. Each entry is an average of two trials unless otherwise indicated.
TON determined by NMR analysis of the production of amide. bOnly one trial.
236
Entry
KOtBu (mol%) TON
1
0 790
2
1 953
3 No catalyst 0 236
4 No catalyst 1 876
Entry
KOtBu (mol%) TON
1
0 463
2
1 860
3 No catalyst 0 563
Table A2.7: Ester Substratea
Table A2.8: Aldehyde Substratea
aReaction conditions: 3 µmol catalyst 1 (0.1 mol%), 3 mmol ester, 12
mmol amine, and X mol% base (potassium tert-butoxide) in 5 mL
THF at 80 °C for 8 hrs. Each entry is an average of two trials unless
otherwise indicated. TON determined by GC analysis of the
production of amide.
aReaction conditions: 3 µmol catalyst 1 (0.1 mol%), 3 mmol aldehyde, 12
mmol amine, and X mol% base (potassium tert-butoxide) in 5 mL THF
at 80 °C for 8 hrs. Each entry is an average of two trials unless otherwise
indicated. TON determined by GC analysis of the consumption of
aldehyde.
237
NMR Studies of the Formation of 2
1
1 + 2 5
B
No reaction
Figure A2.1: Sequential addition of reagents A. Room temperature 1H NMR spectra, hydride region; blue spectrum is the starting
material, red spectrum is after addition of morpholine (no change); green spectrum is
after addition of methanol. B. Room temperature 31P{1H} NMR spectra, same
sample/addition sequence. A small amount of 5 is detected due to catalyst
decomposition. Sample dissolved in benzene-d6.
A
1
1 + 2
No reaction
238
-80 °C Free
13CH
3OH 2 THF THF
-80 °C
22 °C
1+2
A
-80 °C
22 °C
-10 °C
-20 °C
2
3 4 5 Free
Ligand
1+2
B
3 2
1+2
C
-80 °C
22 °C
-10 °C
-20 °C
239
Figure A2.2: Low temperature NMR experiments
Low-temperature NMR experiments with 2, including A. 13
C{1H} NMR
(red and blue) and 13C NMR (green) spectra, B. 31P{1H} NMR spectra, C. 1H NMR spectra, hydride region. The Fe-H peaks for 4 are not observed
due to its low concentration. Sample dissolved in THF-d8.
240
2.3: Iron-Catalyzed Urea Synthesis: Dehydrogenative Coupling of Methanol and
Amines
General Considerations:
All manipulations were carried out under a nitrogen or argon atmosphere using standard
Schlenk, vacuum, cannula, or glovebox techniques. Catalyst 1 was prepared as previously
described.1 Formamides that were not commercially available were prepared using
previously reported procedures.6 All other chemicals were purchased from Aldrich, Fisher,
Strem, Oakwood Chemicals, VWR, or Cambridge Isotope Laboratories. Liquid amine and
alcohol substrates were dried over calcium hydride or sodium hydride, purified by vacuum
transfer or distillation, and stored over 3 Å molecular sieves. Solid substrates were purified
by sublimation, followed by recrystallization (if necessary). Bulk solvents were dried and
deoxygenated using literature procedures.4 NMR solvents were dried over 3 Å molecular
sieves and then used without further manipulation, or sodium and then vacuum transferred
prior to use. Hydrogen was purchased from Airgas and was used as received. 1H, 13C and
31P NMR spectra were recorded on Bruker 300 MHz Avance II+, 300 MHz DRX, 500
MHz DRX or 600 MHz spectrometers at ambient temperature, unless otherwise noted.
Chemical shifts are reported in ppm; J values are given in Hz. 1H and 13C chemical shifts
are referenced to residual solvent signals; 31P chemical shifts are referenced to an external
standard of H3PO4. Probe temperatures were calibrated using ethylene glycol and methanol
as previously described.5 Gas chromatography was performed on a Thermofisher Scientific
Trace 1300 Series gas chromatograph with FID using helium as a carrier gas.
241
General Methods for Symmetric Urea Formation from Catalytic Methanol
Dehydrogenation in the Presence of Amines
In a glovebox, a 100-mL Schlenk tube was loaded with 5 mL of tetrahydrofuran (THF), 12
mmol of amine, 15 µmol of catalyst, and 3 mmol of alcohol, then sealed. It was
immediately placed in an oil bath preheated to 120 °C and stirred for 8 hours. It was then
cooled in an ice bath for 30 minutes. If the starting amine was benzylamine or its
derivatives, all of the following sample preparations and reaction workups were performed
in a glovebox due to the air sensitivity of the remaining starting material.
Analysis of formamide yield:
If analyzed by NMR, 100 µl of reaction solution were added to an NMR tube with 395 µl
CDCl3 and 5 µl of mesitylene standard and an NMR delay time of 60 seconds was used. If
analyzed by GC, 100 µl of reaction solution were diluted to 1 mL with THF and mesitylene
standard (0.024 M or 0.0024M after final dilution) was added.
Urea isolation:
The solvent was then removed from the reaction mixture using glovebox vacuum or rotary
evaporation. The resulting oily solid was transferred to a filtration frit and washed with
room temperature pentane in a glovebox for benzylamine-type starting materials and cold
pentane in air for alkylamine starting materials. The filtrate for alkylamine starting
materials was collected, dried, and re-washed until no more solid was recovered. In the
case of diaminocyclohexane, the product was an oil that was purified with silica gel column
chromatography using 20:1 CH2Cl2:methanol as the eluent according to previously
242
established procedures.7 Ureas were identified by comparison to previously reported
literature data.8-14
General Methods for Unsymmetric Urea Formation from Catalytic Dehydrogenative
Coupling of Formamides and Amines
In a glovebox, a 100-mL Schlenk tube was loaded with 5 mL of tetrahydrofuran (THF), 3
mmol of amine, 15 µmol of catalyst, and 3 mmol of formamide, then sealed. It was
immediately placed in an oil bath preheated to 120 °C and stirred for 16 hours. It was then
cooled in an ice bath for 30 minutes. If the starting amine was benzylamine or its
derivatives, all of the following sample preparations and reaction workups were performed
in a glovebox due to the air sensitivity of the remaining starting material.
Analysis of formamide yield:
If analyzed by NMR, 100 µl of reaction solution were added to an NMR tube with 395 µl
CDCl3 and 5 µl of mesitylene standard and an NMR delay time of 60 seconds was used. If
analyzed by GC, 100 µl of reaction solution were diluted to 1 mL with THF and mesitylene
standard (0.024 M or 0.0024M after final dilution) was added.
Urea isolation:
The solvent was then removed from the reaction mixture using glovebox vacuum or rotary
evaporation. The resulting oily solid was transferred to a filtration frit and washed with
room temperature pentane in a glovebox for benzylamine-type starting materials and cold
pentane in air for alkylamine starting materials. The filtrate for alkylamine starting
materials was collected, dried, and re-washed until no more solid was recovered. In the
243
case of isobutylformamide as the substrate, the product was an oil that was dried under
vacuum and then analyzed by NMR spectroscopy.
244
Optimization Tables for Symmetric Urea Synthesis
Table A2.9: Catalyst Loadinga
Entry Amine Catalyst
Loading
(mol%)
TON (Urea) Yield (%)
1 Cyclohexylamine 0.1 33 3%
2 Cyclohexylamine 0.25 34 9%
3 Cyclohexylamine 0.5 23 12%
Table A2.10: Alcohol:Amine Ratioa
Entry Amine Alcohol:Amine TON (Urea)
1 Cyclohexylamine 1:1 7
2 Cyclohexylamine 1:4 23
3 Cyclohexylamine 1:6 29
4 Cyclohexylamine 1:10 34
Table A2.11: Solventa
Entry Amine Solvent TON (Urea)
1 Cyclohexylamine Ethyl acetate 12
2 Cyclohexylamine 1,4-Dioxane 18
3 Cyclohexylamine Tetrahydrofuran 23
4 Cyclohexylamine Toluene 28
aReaction conditions: X mol% catalyst 1, 3 mmol alcohol, and 12 mmol amine in 5 mL THF at 80°C for
8 hrs. Each entry is an average of two trials unless otherwise indicated. TON determined from isolated
yield of urea.
aReaction conditions: 0.5 mol% catalyst 1, 3 mmol alcohol, and X mmol amine in 5 mL THF at 80°C for
8 hrs. Each entry is an average of two trials unless otherwise indicated. TON determined from isolated
yield of urea. TON error is ±13.
aReaction conditions: 0.5 mol% catalyst 1, 3 mmol alcohol, and 12 mmol amine in 5 mL solvent at 80°C
for 8 hrs. Each entry is an average of two trials unless otherwise indicated. TON determined from isolated
yield of urea.
245
Table A2.12: Temperaturea
Entry Amine Temperature (°C) TON (Urea)
1 Cyclohexylamine 60b 5
2 Cyclohexylamine 80 23
3 Cyclohexylamine 80c 28
4 Cyclohexylamine 100 55
5d Cyclohexylamine 120 66
6 Cyclohexylamine 120c 66
Table A2.13: Timea
Entry Amine Time (hrs) TON (Urea)
1 Cyclohexylamine 1 11
2b Cyclohexylamine 4 44
3b Cyclohexylamine 8 66
4b Cyclohexylamine 16 56
Table A2.14: Addition of Basea
Entry Amine DBU Additive
(mol%)
TON (Urea)
1 4-(trifluoromethyl)
benzylamine
0 90
2b 4-(trifluoromethyl)
benzylamine
5 108
aReaction conditions: 0.5 mol% catalyst 1, 3 mmol alcohol, and 12 mmol amine in 5 mL THF at X°C for
8 hrs. Each entry is an average of two trials unless otherwise indicated. TON determined from isolated
yield of urea. b16 hrs. cToluene solvent. dAverage of three trials.
aReaction conditions: 0.5 mol% catalyst 1, 3 mmol alcohol, and 12 mmol amine in 5 mL THF at 120°C
for X hrs. Each entry is an average of two trials unless otherwise indicated. TON determined from isolated
yield of urea. bAverage of three trials.
aReaction conditions: 0.5 mol% catalyst 1, 3 mmol alcohol, and 12 mmol amine in 5 mL THF at 120°C
for 8 hrs, with X mol% 1,8-diazabicyclo(5.4.0)undec-7-ene (DBU). Each entry is an average of three
trials unless otherwise indicated. TON determined from isolated yield of urea. bOnly one trial.
246
Optimization Tables for Unsymmetric Urea Synthesis
Table A2.15: Timea
Entry R R’ Time (hrs) TON (Urea)
1 Benzyl Cyclohexyl 4b 0
2 Benzyl Cyclohexyl 4 68
3 Benzyl Cyclohexyl 8 156
4 Benzyl Cyclohexyl 16 170
Table A2.16: Temperaturea
Entry R R’ Temperature
(°C)
TON (Urea)
1 Benzyl Cyclohexyl 100 162
2 Benzyl Cyclohexyl 120 170
3 Benzyl Cyclohexyl 140 138
Table A2.17: Catalyst Loadinga
Entry R R’ Catalyst
Loading
(mol%)
TON
(Urea)
Yield (%)
1 Benzyl Cyclohexyl 0.5 170 85%
2 Benzyl Cyclohexyl 1 83 83%
aReaction conditions: 0.5 mol% catalyst 1, 3 mmol formamide, and 3 mmol amine in 5 mL THF at 120°C
for X hrs. Each entry is an average of two trials unless otherwise indicated. TON determined from isolated
yield of unsymmetric urea. bNo catalyst.
aReaction conditions: 0.5 mol% catalyst 1, 3 mmol formamide, and 3 mmol amine in 5 mL THF at X°C
for 16 hrs. Each entry is an average of two trials unless otherwise indicated. TON determined from
isolated yield of unsymmetric urea.
aReaction conditions: X mol% catalyst 1, 3 mmol formamide, and 3 mmol amine in 5 mL THF at 120°C
for 16 hrs. Each entry is an average of two trials unless otherwise indicated. TON determined from
isolated yield of unsymmetric urea.
247
Table A2.18: Formamide Scramblinga
Entry R R’ Starting
Formamide
Left
(mmol)
Scrambled
Formamide
(mmol)
Catalyst Yield,
Unsymm.
Urea (%)
Sel.b
(%)
1 Cyclo-
hexyl
Benzyl 1.53 0.02 1, 0.5
mol%
34% 72%
2 Cyclo-
hexyl
Benzyl 2.70 0.04 None 0 NA
aReaction conditions: 0.5 mol% catalyst 1 or no catalyst, 3 mmol formamide, and 3 mmol amine in 5 mL
THF at 120°C for 4 hrs. Each entry represents only one trial. NA = not applicable. bSelectivity determined
by amount of unsymmetric urea as compared to the amounts of the two symmetric ureas.
248
Mechanistic NMR Experiments
Figure A2.3: Hydrogenation of ureas
Room temperature 13C{1H} NMR spectra in THF-d8, 600 MHz, carbonyl region; blue spectrum is straight
after mixing at room temperature, red spectrum is after 16 hrs. at 120 °C. GC-FID showed only ~1.5%
conversion to formamide.
Immediate
16 hrs. @ 120 °C
1,3-dibenzylurea
249
Figure A2.4: Reaction of urea with amine under H2
Room temperature 13C{1H} NMR spectrum in THF-d8, 600 MHz, carbonyl region; after 16 hrs. at 120°C.
The unsymmetric urea product is marked by a box.
1,3-dibenzylurea
Unsymmetric Urea
250
Figure A2.5: Reaction of urea with formamide under H2
Room temperature 13C{1H} NMR spectrum in THF-d8, 600 MHz, carbonyl region; after 16 hrs. at 120
°C. The unsymmetric urea product is marked by a box. The peak at 160.6 ppm is a second rotamer of the
cyclohexylformamide reactant.
Cyclohexylformamide
1,3-dibenzylurea
Unsymmetric Urea
251
Figure A2.6: Reaction of urea with formamide, no H2 Room temperature 13C{1H} NMR spectrum in THF-d8, 600 MHz, carbonyl region; after 16 hrs. at 120
°C. The unsymmetric urea product is marked by a box. The peak at 160.6 ppm is a second rotamer of the
cyclohexylformamide reactant.
1,3-dibenzylurea
Cyclohexylformamide
Unsymmetric Urea
252
Figure A2.7: Reaction of formamides and amines Room temperature 13C{1H} NMR spectra in THF-d8, 600 MHz, carbonyl region; blue spectrum is straight
after mixing at room temperature, red spectrum is after 2 hrs. at 120 °C, green spectrum is after 5 hrs.
Immediate
2 hrs. @ 120 °C
5 hrs. @ 120 °C
Unsymmetric Urea
Benzylformamide
Cyclohexyl
formamide
253
Table A2.19: Attempt to Make a Tetrasubstituted Urea: Standard Catalytic
Conditionsa
Entry Formamide Left Urea Produced
1 3 mmol 0
Immediate
16 hrs. @ 120 °C
Figure A2.8: Attempt to make a tetrasubstituted urea: NMR-scale
Room temperature 13C{1H} NMR spectra in THF-d8, 600 MHz carbonyl region; blue spectrum is straight
after mixing at room temperature, red spectrum is after 16 hrs. at 120 °C.
No reaction
aReaction conditions: 0.5 mol% catalyst 1, 3 mmol formamide, and 12 mmol amine in 5 mL THF at 80°C
for 8 hrs. Each entry represents only one trial.
No reaction
254
Figure A2.9: Reaction of pentylformamide with dipentylamine
Room temperature 13C{1H} NMR spectra in THF-d8, 600 MHz, carbonyl region; blue spectrum is straight
after mixing at room temperature, red spectrum is after 1 hr. at 120 °C, green spectrum is after 2 hrs., and
purple spectrum is after 24 hrs. Carbonyl carbon for the trisubstiuted urea is shown by the blue box. The
other carbonyl carbon is from the starting formamide. GC-FID showed presence of the trisubstituted urea
and no dipentylformamide or pentylamine from scrambling.
Immediate
24 hrs. @ 120 °C
2 hrs. @ 120 °C
1 hr. @ 120 °C
Unsymmetric Urea
255
No reaction
Immediate
1 hr. @ 120 °C
2 hrs. @ 120 °C
24 hrs. @ 120 °C
Figure A2.10: Reaction of dipentylformamide with pentylamine
Room temperature 13C{1H} NMR spectra in THF-d8, 600 MHz, carbonyl region; blue spectrum is straight
after mixing at room temperature, red spectrum is after 1 hr. at 120 °C, green spectrum is after 2 hrs., and
purple spectrum is after 24 hrs. The carbonyl carbon shown is from the starting formamide. GC-FID
showed no trisubstituted urea and no pentylformamide or dipentylamine from scrambling.
256
A.
2 1
3
257
B.
Figure A2.11: Dehydrogenation of a formamide to an isocyanate
A. Room temperature 1H NMR spectrum in THF-d8, hydride region, after 1 hr. at 120 °C. Presence of the
iron-dihydride (trans isomer) is marked by a box. B. 31P{1H} NMR spectrum in THF-d8 after 1 hr. at 120
°C. Presence of 5 and some free ligand due to catalyst decomposition. Arrows indicate phosphorous
resonances for both the cis (red) and trans (green) dihydride isomers. 300+ MHz. GC-FID showed
presence of cyclohexyl isocyanate.
2
1
Free
Ligand
5
4 3
258
NMR Spectra for Synthesized Symmetric Ureas
All 1H NMRs are taken on a 300+ MHz spectrometer and all 13C{1H} NMRs are taken on
a 600 MHz spectrometer unless otherwise indicated. * = small amount of formamide
impurity
259
1,3-di-n-pentylurea (Table 2.3.1, Entry 1)
1H NMR, CDCl3
13C{1H} NMR, CDCl3
*
*
260
1,3-bis(2-methylpropyl)urea (Table 2.3.1, Entry 2)
1H NMR, CDCl3
13C{1H} NMR, CDCl3
*
*
261
1,3-di-n-heptylurea (Table 2.3.1, Entry 3)
1H NMR, CDCl3
13C{1H} NMR, CDCl3
262
1,3-bis(2-methoxyethyl)urea (Table 2.3.1, Entry 4)
1H NMR, CDCl3
13C{1H} NMR, CDCl3
*
*
263
1,3-bis(1-methylhexyl)urea (Table 2.3.1, Entry 5)
1H NMR, CDCl3
13C{1H} NMR, CDCl3
264
1,3-bis(4-methoxybenzyl)urea (Table 2.3.1, Entry 6)
1H NMR, CDCl3
13C{1H} NMR, CDCl3
*
*
265
1,3-bis(4-methylbenzyl)urea (Table 2.3.1, Entry 7)
1H NMR, CDCl3
13C{1H} NMR, CDCl3
266
1,3-dibenzylurea (Table 2.3.1, Entry 8)
1H NMR, CDCl3
13C{1H} NMR, CDCl3
*
*
267
1,3-bis(4-trifluoromethylbenzyl)urea (Table 2.3.1, Entry 9)
1H NMR, CDCl3
13C{1H} NMR, CDCl3
*
268
1,3-bis(2-phenylethyl)urea (Table 2.3.1, Entry 10)
1H NMR, CDCl3
13C{1H} NMR, CDCl3
269
1,3-dicyclopentylurea (Table 2.3.1, Entry 11)
1H NMR, CDCl3
13C{1H} NMR, CDCl3
270
1,3-dicyclohexylurea (Table 2.3.1, Entry 12)
1H NMR, CDCl3
13C{1H} NMR, CDCl3
271
Octahydro-benzoimidazol-2-one (Table 2.3.1, Entry 13)
Note: observe both cis and trans isomers of the product urea
1H NMR, CDCl3
13C{1H} NMR, CDCl3
272
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