Homogeneous Catalysis for Sustainable Energy: Hydrogen
and Methanol Economies, Fuels from Biomass, and Related Topics
As the world pledges to significantly cut carbon emissions, the demand for sustainable and clean energy has now become more important than ever. This includes both production and storage of energy carriers, a majority of which involve catalytic reactions. This article reviews recent developments of homogeneous catalysts in emerging applications of sustainable energy. The most important focus has been on hydrogen storage as several efficient homogeneous catalysts have been reported recently for (de)hydrogenative transformations promising to the hydrogen economy. Another direction that has been extensively covered in this review is that of the methanol economy. Homogeneous catalysts investigated for the production of methanol from CO2, CO, and HCOOH have been discussed in detail. Moreover, catalytic processes for the production of conventional fuels (higher alkanes such as diesel, wax) from biomass or lower alkanes have also been discussed. A section has also been dedicated to the production of ethylene glycol from CO and H2 using homogeneous catalysts. Well-defined transition metal complexes, in particular, pincer complexes, have been discussed in more detail due to their high activity and well-studied mechanisms.
Introduction<!>General Mechanistic Consideration<!>General Classes of Catalysts Discussed in
This Review<!>General Mechanistic Consideration<!>Hydrogen Economy<!>Methanol
as a Hydrogen Storage Material<!>Aqueous Methanol Reforming Using Ruthenium-Based
Catalysts<!>Aqueous Methanol Reforming
Using Iridium
and Rhodium-Based Catalysts<!>Aqueous Methanol Reforming Using Base-Metal
Catalysts<!>Hydrogen Production from
Aqueous Formaldehyde/Paraformaldehyde<!>Dehydrogenation of Aqueous Formaldehyde/Paraformaldehyde<!>Hydrogen Production from
Aqueous Formaldehyde/Paraformaldehyde<!>Interconversion of
Complexes Ir4, Ir4a,
and Ir4b on Changing the pH<!>Hydrogen Production from
Aqueous Formaldehyde/Paraformaldehyde<!>Synthesis of Complexes Ru3ka, Ru3kb, and Ru3kc and Catalytic
Dehydrogenation of Aqueous HCHO<!>Hydrogen Production from
Aqueous Formaldehyde/Paraformaldehyde<!>Amine-boranes
as Hydrogen Storage Materials<!>Dehydrogenation of Linear Amine-boranes<!>General Pathways for the Dehydrocoupling of Linear
Amine-boranes159<!>Dehydrogenation
of Cyclic Amine-boranes<!>Regeneration
of Amine-boranes from the Spent
Fuel<!>Liquid Organic Hydrogen Carriers (LOHCs)<!><!>LOHCs Based on Carbocycles<!>LOHCs Based on Heterocycles<!>LOHCs
Based on Alcohols and Amines<!>Proposed Mechanism for the Dehydrogenative Coupling
of DMEDA and
Methanol<!>LOHCs
Based on Alcohols and Amines<!>LOHC Based on 1,4-Butanediol and γ-Butyrolactone<!>LOHCs
Based on Alcohols and Amines<!>Proposed Catalytic
Cycle for Dehydrogenative Coupling of EG to Form
Polyester Using a Ruthenium Catalyst<!>LOHCs
Based on Formic Acid and CO2<!><!>LOHCs
Based on Formic Acid and CO2<!>Proposed Mechanism of the Hydrogenation of CO2 to Formate
Using a Ru Pincer Catalyst<!>Hydrogen Production from Biomass and Water
Splitting<!>Advanced
Biofuel from Ethanol<!>Guerbet Reaction Catalyzed by Homogeneous (Nonpincer) Catalysts<!>Advanced
Biofuel from Ethanol<!>Hydrogenation
of Vegetable Oils: Upgrading
Biodiesel<!>Lignin
Depolymerization<!>Polymeric
Structure of Lignin<!>Lignin
Depolymerization<!>Methanol
Economy<!>Methanol
Production from CO2<!>CO2 to Methanol
Using Amine as
a Capturing Agent<!>Proposed Mechanism
of Amine-Assisted CO2-to-Methanol Transformation
Using Ru Pincer Complex<!>CO2 to Methanol
Using Amine as
a Capturing Agent<!>CO2 to Methanol Using Alcohols<!>CO2 to Methanol Using Silanes/Boranes<!>Methanol Production from CO<!>Methanol Production from Formic Acid<!>Formic Acid Decomposition to Methanol<!>Methanol Production from Methane<!>Alkane Upgrading to Liquid Fuels<!>Alkane
Metathesis<!>Cross Alkane Metathesis Involving (De)hydrogenation
and Olefin Metathesis416<!>Alkane
Metathesis<!>Alkane–Alkene Coupling<!>Ethylene Glycol Production from CO and H2<!>Summary and Outlook<!>Author Contributions<!>
<p>Energy lies at the core of a nation's economy, and for the
past two centuries, the amount of the energy consumption of society
has increased in lockstep with the amount of wealth created. Previous
data indicate a correlation between energy consumption and gross domestic
product (GDP) implicating increased demand for energy globally, especially
for the developing economy.1 On the basis
of the recent data, the world's primary energy source constitutes
petroleum (34%), coal (27%), and natural gas (24%), making 85% of
the energy sources to be fossil fuels.2 Heavy consumption of fossil fuels is not sustainable as they take
millions of years to form and the limited supply of fossil fuels is
being depleted at a much faster rate than they are being produced.
Additionally, the consumption of fossil fuels raises serious environmental
and health concerns. For example, a vast majority of deaths due to
air pollution are caused due to fossil fuel consumption and more than
3 million lives would be saved every year if cleaner energy were used
instead.3 Unsurprisingly, a direct correlation
between GDP and the amount of CO2 being produced also exists.4,5 Thus, there is an urgent need to develop sustainable and clean energy
carriers.</p><p>Several alternate energy sources such as solar, wind,
ocean (tidal,
wave, thermal), biomass, nuclear, and geothermal have been well-studied
in the past, where each one has its own limitations, and for a practical
scenario in the near future, a combination of multiple renewable energy
sources together with fossil fuels will be needed to keep the planet
sustainable and green.6−8 Pursuit of a suitable carrier for the production,
storage, and use of energy in a clean, economical, and sustainable
way has created the need for the development of efficient catalysts
to advance this goal. In general, catalytic technologies play crucial
roles in the following sections of the energy sector: (a) energy production
reactions, (b) safe and long-term energy storage and transportation,
and (c) high efficiency of energy use. Multiple review articles have
been reported recently on the application of heterogeneous catalysts,9−13 photocatalysts,14,15 and electrocatalysts16−21 for the production and storage of energy, and these subdisciplines
of catalysis will not be discussed here.</p><p>Homogeneous catalysis
allows the processes to occur under relatively
mild conditions and at the same time advances our understanding of
reaction mechanisms at the molecular level, thus providing remarkable
opportunities to improve the catalytic processes.22 Here, we review reports on homogeneous catalysis based
on transition metal complexes for their applications in the development
of clean and sustainable energy carriers. Emphasis has been given
to complexes based on pincer23 type ligands
as they have led this area and let some remarkable discoveries happen.
Pincer ligands are defined as chelating ligands that bind through
three adjacent donor sites in a meridional geometry.24 The choice of topics as detailed below is based on energy
production or storage systems in which homogeneous transition metal
catalysis has played a prominent role.</p><p>H2, long considered
a "fuel of the future",
in the past decade has climbed its way up to the stage where a hydrogen
economy25 is looking promising in the foreseeable
future. Several well-defined transition metal catalysts have been
investigated in the past to impact the hydrogen economy, especially
to discover new hydrogen storage systems. Section 2 of this review discusses various hydrogen
storage systems and the development of metal complexes as catalysts
for the charge and discharge of H2. Aqueous reforming of
methanol (CH3OH + H2O = 3H2 + CO2) using both precious metal and earth abundant metal based
catalysts have been discussed in detail because of the potential application
of this reaction to impact both the hydrogen economy and the methanol
economy (section 2.1).26 Recent developments in the direction
of production of H2 from formaldehyde/paraformaldehyde
have also been discussed (section 2.2). The reaction is catalyzed by ruthenium or iridium
complexes and produces CO2 or carbonate salts as byproducts.
Of many inorganic materials explored for the purpose of hydrogen storage,
amine-boranes have been the most studied by the community of organometallic
chemists. Catalytic dehydrogenation of amine-boranes has only been
briefly discussed here as multiple review articles have been reported
in the past few years on this topic (section 2.3).27−36 Despite significant developments in the dehydrogenation of various
amine-boranes, a practical technology for the regeneration of "charged
fuel" (amine-boranes) from "spent fuel" (e.g.,
borazine) is yet to be developed. A perspective on the regeneration
of amine-borane "charged fuel" and developments in the
area of hybrid cyclic amine-boranes has also been detailed (section 2.3.3). The quest
of developing reversible hydrogen storage materials has led to the
development of "liquid organic hydrogen carriers" (LOHCs).
Various LOHCs have been developed in the past; some of them have also
been commercialized.37 Recent advances
in the development of various LOHCs using organometallic catalysts
have been discussed in detail (section 2.4). A brief section on the production of
H2 from biomass and water splitting has also been included
here (section 2.5).</p><p>The use of renewable feedstock or biomass to produce valuable
chemicals
and energy carriers lies at the heart of the circular economy model.38−40 Fuels produced from biomass or "biofuels" can be categorized
into three types. The first type, called the "first-generation
biofuels", mainly involving bioethanol and biodiesel, is produced
from edible biomass such as sugars, grains, or seeds.41 Both bioethanol and biodiesels can be used to substitute
petrol and diesel respectively, most commonly as blends with conventional
fuels. Bioethanol is produced through the fermentation of sugar via
enzymatic catalysis,42,43 whereas biodiesel is produced
through the transesterification of vegetable oils or fats for which
several types of catalysts, e.g., acid/base catalysts,44 enzymes,45 and heterogeneous
catalysts,46,47 have been reported. Recent advances
have demonstrated that ethanol can be converted to butanol and other
higher alcohols which have multiple advantages (e.g., high energy
density, noncorrosive nature, and immiscibility with water) over ethanol,
making them closer to the conventional gasoline fuels. We have reviewed
here recent progress made in the direction of transforming ethanol
to butanol and higher alcohols using organometallic catalysts (section 3.1). Additionally,
we have also discussed reports on the catalytic hydrogenation of the
unsaturated C=C bonds in the fatty acid methyl esters to upgrade
biodiesels by improving their multiple properties, e.g., stability,
and lubricity (section 3.2). Technologies for the production of first-generation biofuels
are well-established and are operating around the globe. However,
it is important to note that the first-generation biofuels are produced
from edible food crops that make their production process compete
with the land and water used for food and it has been claimed that
they have driven up the cost of food and animal feeds.48,49</p><p>The second type of biofuel, called "second-generation
biofuels",
is sourced from nonedible biomass, e.g., lignocellulose (e.g., cereal
straw, sugar cane bagasse, and organic waste). This avoids the complexity
of the societal implication of using food to produce fuels. Technologies
to transform lignocellulose to fuels, especially alternative jet fuels,
are being evaluated.50,51 These processes involve catalytic
pyrolysis or gasification and Fischer–Tropsch reactions followed
by hydrotreatment.52,53 The main limitation of "second-generation
biofuel" is the requirement of sophisticated processing and
production equipment as the complex 3-D structure of lignocellulose
makes it difficult to depolymerize. Several heterogeneous and acid/base
catalysts have been used to depolymerize lignin under harsh conditions
as reported in the past.54−62 Homogeneous transition metal catalysis has also been employed to
break model compounds of lignin as discussed in this review (section 3.3).</p><p>The
third type of biofuel, called "third-generation biofuels",
involves fuels derived from microorganisms, e.g., algae. Despite some
important advantages such as high ignition points, and biodegradability,
the high cost associated with the production of fuels from algae has
limited its scope for commercial purposes.63</p><p>Another area that has attracted significant interest in the
past
decade is that of the methanol economy. Methanol is an important chemical
feedstock for plastics, glue, paints, building materials, and solvents
and is produced on a scale of more than 75 million tons annually worldwide.
Its additional application as a fuel (or fuel additive) has sparked
further interest in this area. A sustainable, cost-effective, and
large-scale production of methanol from captured CO2 would
be highly useful to enable a "circular economy"38 as also proposed by Olah and Prakash as the
concept of "methanol economy".64 Beyond the concept, this vision has been realized by an Iceland-based
company called Carbon Recycling International (CRI), where the CO2 captured from industrial emissions is hydrogenated to methanol
using renewable H2. The renewable H2 can come
from the electrolysis of water using renewable electricity (making
renewable e-methanol or Vulcanol) or from the byproduct or waste gas
(making low carbon methanol). Since 2021, Vulcanol has been commercially
sold in Europe and China, and CRI aims to produce up to 110 000
tons of recycled carbon methanol annually from 2021.65 Another company that makes renewable methanol on an industrial
scale is the Canada-based company called Enerkem,66 where methanol is produced from municipal solid waste via
thermochemical gasification of organic waste to produce syngas followed
by the catalytic conversion of syngas to methanol. The Netherland-based
company called BioMCN67 produces renewable
methanol from biogas sourced from waste digestion plants. In 2017,
BioMCN produced and sold around 60 000 tons of renewable methanol.
It is clear that the production of renewable methanol is still much
lower than its demand, which is partly due to the lack of an abundant
supply of inexpensive renewable feedstocks or renewable energy in
the local region. Considering that the production and release of CO2 in the atmosphere is inevitable, its capture and transformation
to methanol in a cost-effective manner is perhaps the most sustainable
approach to produce renewable methanol. This area has been extensively
studied using organometallic catalysts where CO2 can be
trapped using capturing agents such as alcohols, amines, silanes,
and boranes followed by its hydrogenation or hydrolysis to produce
methanol. We have reviewed here examples of the indirect transformation
of CO2 to methanol using homogeneous catalysts (section 4.1). A brief discussion
on the production of methanol from CO (section 4.2), HCOOH (section 4.3), and CH4 (section 4.4) catalyzed by transition
metal complexes has also been reported.</p><p>The above-described
topics investigate the pursuit of alternative
energy sources such as H2, methanol, and biofuels. Another
approach to manifest sustainable energy can be to develop methods
to convert renewable feedstock to conventional fossil fuels. For example,
lower alkanes, e.g., methane, can be produced from biomass or CO2. Thus, its transformation to produce higher hydrocarbons
(e.g., C8–C19) can allow us to attain a sustainable energy-based
economy without a need to change our infrastructure to accommodate
a new energy source. This is particularly important for aviation fuels,
where density requirements and stringent specifications dictate that
jet fuel (C8–C16 hydrocarbons) will be the industry norm in
the near future.68 The area of alkane upgradation
has made slow progress in the past couple of decades due to the inertness
and inactivity of alkane C–H bonds. Two approaches using homogeneous
catalysts have been utilized for alkane upgradation and demonstrate
promise in this direction. We have reviewed here these two approaches
that are based on alkane metathesis (section 5.1) and alkane–alkene coupling (section 5.2). A brief section
on the use of ethylene glycol as a fuel and its production using homogeneous
catalysts has also been presented (section 6). Finally, a summary and perspective on
the current challenges and prospects of the reviewed areas have been
described (section 7).</p><!><p>A drive
to study homogeneous catalysis is the opportunity to understand
the mechanisms of catalytic processes which eventually allows us to
develop new and more efficient catalysts. As we review a plethora
of homogeneous catalysts for the production and storage of energy
carriers, in this section, we present a general overview of their
mechanism of operation. Many such catalysts operate via redox innocence
metal–ligand cooperation (MLC) featuring a basic or nucleophilic
site on the ligand, and a coordinatively unsaturated metal center
that can act as an electrophile. This allows heterolytic bond activation
of polar (e.g., O–H, N–H) and nonpolar (e.g., H2) molecules across the electrophilic metal center and the
nucleophilic ligand site. Overall, the oxidation state of the metal
remains the same, unlike the classical mode of bond activation by
oxidative addition and reductive elimination, where the oxidation
state of the metal changes by two units. This has allowed the utilization
of this concept in bond activation and catalysis to various elements
(including main-group elements such as boron69 and zinc70) across the periodic table.
Overall, most of the catalysts discussed in this article can be broadly
categorized into the following types depending on their mode of operation.</p><p>(a) MLC via amido/amino mode. Application of MLC via
amido/amino mode was first demonstrated by Noyori for catalytic hydrogenation
reactions. Since then several examples of transition metal complexes
exhibiting MLC via amido/amino have been studied as reviewed in the
recent past.71 An example of this mode
operation is shown in Scheme 1A using a pincer complex containing a MACHO-type ligand. Ruthenium
MACHO complexes exhibit tolerance and robustness toward harsh catalytic
conditions such as temperature (up to 150 °C), because of which
they have been extensively employed for a variety of green homogeneous
catalysis, in particular (de)hydrogenative transformations. Examples
of such catalysts can be seen throughout this review using complexes
of Ru, Ir, Mn, and Fe.</p><!><p>X and Y are neutral ligands,
e.g., phosphine or amine derivatives. HA represents bonds such as
H–H, H–C, H–O, H–N, H–B, and H–Si.
R = tBu, Ph; R′ = benzyl, tBu, and iPr.
D = neutral two-electron donor site.</p><!><p>(b) MLC via dearomatization/aromatization mode. A
new mode of metal–ligand cooperation based on dearomatization
and aromatization of pincer complexes was reported by Milstein. As
shown in Scheme 1A,
such complexes involve pincer ligands containing lutidine backbone
and bond activation occurs between the side arm CH proton and the
metal center. The dearomatization step is driven by the formation
of a stronger bond between the formed amido ligand and metal, whereas
the aromatization process is driven by stability gained due to aromatization
of the pyridine ring and formation of a more stable coordinatively
saturated complex. Extensive investigations by both computation and
experiments have been carried to understand the mechanism in detail.
A recent article by Khaskin and Gusev reports that the bipyridine
complex Ru-BipyPNN can undergo hydrogenation of the bipyridine central
ring under reducing conditions.72 A review
on the application of catalysts exhibiting metal–ligand cooperation
via dearomatization/aromatization mode has recently been reported
by Fujita.73 Another related complex is
based on acridine-type PNP ligands where a dearomatized acridine complex
is found to be active in catalysis as discussed in this review (e.g., Scheme 24).</p><p>(c) Dual mode of MLC. The Milstein group has recently
reported a new family of pincer complexes capable of exhibiting a
dual mode of MLC via an amido/amino mode and a dearomatization/aromatization
mode (Scheme 1A).74 Such catalysts of ruthenium are highly active
for (de)hydrogenative catalysis and can hydrogenate esters and amides75 near room temperature. Such catalysts have also
been applied for the hydrogenation of commercial resins of nylons
and polyurethanes.76</p><p>(d) Redox-active
MLC. Grützmacher has developed
metal complexes containing the 1,4-bis(5H-dibenzo[a,d]cyclohepten-5-yl)-1,4-diazabuta-1,3-diene
ligand that have demonstrated promising activities for the production
of H2 from aqueous methanol77 or formaldehyde78 as discussed in this
review. Such complexes exhibit both redox noninnocence and chemical
noninnocence as shown in Scheme 1B.</p><p>(e) Catalysts based
on proton responsive ligands. Complexes
bearing proton responsive ligands that can be deprotonated and protonated
on change of pH have been utilized for (de)hydrogenative transformation
reactions by multiple research groups (Scheme 1C).79,80 The deprotonated and
protonated complexes can catalyze different reactions. For example,
the deprotonated catalyst can catalyze the hydrogenation of CO2 to formate at higher pH whereas the protonated catalyst can
catalyze the dehydrogenation of HCOOH at a lower pH.80 Such catalysts have been utilized for several (de)hydrogenation
reactions such as aqueous methanol reforming (section 2.1.2), dehydrogenation of
aqueous formaldehyde (section 2.2), and dehydrogenation of HCOOH and the reverse reaction—hydrogenation
of CO2 (section 2.4.4).</p><p>(f) Catalysts involving triphos ligands. Another important
class of catalysts employed in (de)hydrogenative catalysis involves
triphos ligands (e.g., Scheme 53, vide infra). These catalysts are
activated using an acid such as HNTf2, unlike catalysts
exhibiting MLC via the amido/amino or dearomatization/aromatization
mode that operate under either basic or neutral conditions. This complements
the two types of catalysts to work under different pH conditions and
tolerate functional groups of acidic or basic nature.</p><!><p>Hydrogen has been long termed as the
ideal energy source of the
future. Molecular H2 is light, is storable, has the highest
gravimetric energy content of common fuels (120 MJ kg–1), and does not produce any direct emission of common pollutants
or greenhouse gases, making it a very attractive candidate as a sustainable
and clean energy carrier.25,81,82 However, as of now, H2 is primarily used as an industrial
feedstock for the production of chemicals, e.g., ammonia, methanol,
and petroleum refining.83 Application of
(de)hydrogenation reactions to convert waste to useful chemical resources
to enable a circular economy has been recently reviewed.84 Cost-effective and sustainable demonstration
of H2 as a clean energy carrier to manifest the hydrogen
economy faces two major challenges: (a) sustainable production of
renewable H2 and (b) efficient storage of H2.</p><p>More than 95% of H2 is currently produced from
fossil
fuel via steam reforming, coal gasification, or the steam methane
reforming (SMR) process where natural gas (primarily CH4) is reacted with steam to produce CO and H2 (CH4 + H2O → CO + 3H2).85 The produced CO subsequently reacts with steam to produce
more H2 and CO2; the process is known as the
water gas shift reaction (CO + H2O → CO2 + H2).86 Thus, one molecule
of CO2 is produced per four molecules of H2 produced
via the SMR process. This makes the release of 5.5 tons of CO2 in the atmosphere per ton of H2 produced. It is
possible to capture almost 70% of the released CO2 and
store it in deep underground wells (carbon capture and sequestration);87 however, such operations are not in common practice.
Thus, the production of H2 from SMR is not ideal for the
hydrogen economy as the feedstock to produce H2 is a fossil
fuel and a significant amount of CO2 is emitted to the
atmosphere in this process. Therefore, the development of alternative
technologies for the cleaner production of renewable H2 is crucial for the hydrogen economy. Several alternative technologies
for the production of clean and renewable H2 have been
evaluated. A technology that is being tested on a large scale for
the clean production of H2 is the pyrolysis of methane,88 where methane is bubbled on a molten catalyst
at high temperature (∼1000 °C). The reaction (CH4 → C + 2H2; ΔH° = 74
kJ/mol) requires 5 kWh of electricity for process heat to produce
1 kg of H2. The process is claimed to generate no pollution
as the produced carbon can be used as manufacturing feedstock in industry
or be landfilled. Along a similar direction, the Kværner process89 has been developed by the Norwegian engineering
firm Kværner that produces H2 from hydrocarbons (e.g.,
natural gas and biogas) in a plasma burner at 1600 °C (CnHm → nC + m/2H2). Although the production
of H2 from methane pyrolysis and the Kværner process
are clean as they do not produce any greenhouse gas, the produced
H2 is nonrenewable due to the use of fossil fuel feedstock.
Biomass has been investigated as an attractive feedstock to produce
renewable H2 via the process of pyrolysis (biomass + heat
→ CH4 + CO + CO2 + H2 + other
products) or gasification (biomass + O2 or H2O + heat → H2O + CO + CO2 + CH4 + H2 + other products).90 However,
these processes produce CO2, making them less ideal for
the vision of hydrogen economy. The most promising approach for the
clean production of renewable H2 is the electrolysis of
water. However, water electrolysis is thermodynamically unfavorable
(H2O → H2 + 1/2O2; ΔG° = 237.24 kJ/mol; ΔH°
= 285.83 kJ/mol) and requires a substantial energy input (50–55
kWh of electricity/kg of H2).91 This makes the success of the technology dependent on the cost and
nature of the electricity (e.g., renewable or nonrenewable) to be
used for the process. Several renewable energy sources, e.g., wind,
solar, and geothermal energy, have been evaluated or demonstrated
for the large-scale electrolysis of water. In 2019, only 0.1% of the
global hydrogen was produced via the electrolysis of water.92 A brief report on the production of H2 from biomass and water electrolysis using homogeneous transition
metal catalysts is discussed in section 2.5.</p><p>Another major challenge lying in
front of the hydrogen economy
is hydrogen storage as H2 has a very low volumetric energy
density (0.0108 MJ L–1), making it almost impossible
to use in its normal form under mild conditions of pressure and temperature
for several sectors such as transport. Conventionally, hydrogen gas
is stored physically in compressed form at very high pressure (100–700
bar) or in the cryogenic form at a very low temperature of −253
°C. Both the processes are highly energy intensive and not economical,
especially for long-term or long-distance transport. Thus, the safe
and economical storage of hydrogen gas for long-term and long-distance
transport is an important challenge in front of the hydrogen economy.
In recent years, there has been significant development toward different
approaches for hydrogen storage, in particular, by physisorption in
porous materials or by making or breaking chemical bonds. Several
reviews have been reported discussing the advantages and disadvantages
of each technique.93−97 Several properties need to be considered for a suitable hydrogen
storage material, for example, the following properties.</p><p>(a) Gravimetric storage capacity. The material should
have a high gravimetric hydrogen storage capacity (e.g., >5.5 wt
%,
U.S. Department of Energy target).98</p><p>(b) Viscosity. Lower viscosity would allow smoother
transport of the carrier materials through various parts of the reactor/storage
system, thus making a better hydrogen storage material.</p><p>(c) Gas stream purity. The presence of contaminants
such as CO or NH3 in the produced hydrogen gas stream could
poison catalysts used in the fuel cells and reduce the efficiency
of energy production. Thus, materials or processes producing contaminant-free
hydrogen gas are highly desirable.</p><p>(d) Reversibility. For the hydrogen storage process
to be sustainable, both the discharge of H2 from the charged
fuel and the regeneration of the charged fuel from spent fuel need
to be feasible in a green and cost-effective way. Thus, the thermodynamics
of both the dehydrogenation and regeneration processes need to be
considered for the purpose of hydrogen storage.</p><p>(e) Stability. Both the charged fuel and spent fuel
should be thermally/photolytically stable to keep the charging/discharging
cycle continuing.</p><p>(f) Toxicity. The material should
be of low toxicity.</p><p>(g) Availability. The material
should be inexpensive
and abundant. Ideally, it should be compatible with the existing infrastructure.</p><p>H2 gas can be stored through physical adsorption on
various materials. However, in almost all cases either a very low
temperature or a high pressure is required for the storage. Another
area of hydrogen storage is based on chemical hydrogen storage materials
where hydrogen atoms are covalently bound, and the H2 can
be produced by a thermal or catalytic process. Several solid carriers
such as metal hydrides,99,100 borohydrides,101 alanates,102 and
imides/amides103 have been investigated.
A majority of them suffer from issues of low hydrogen storage capacity
and difficulty in the regeneration of the charged fuel from the spent
fuel. Another mode of hydrogen storage is based on liquid organic
hydrogen carriers (LOHCs). Effective release of H2 from
the charged fuel carrier and its selective regeneration under mild
conditions require a suitable catalyst. In this section, we review
how homogeneous catalysis has contributed to the discoveries of new
and potential chemical hydrogen storage materials.</p><!><p>Methanol is an inexpensive
alcohol, is a stable liquid with low
viscosity (∼0.54 mPa·s) under ambient conditions, and
has a high theoretical gravimetric hydrogen capacity (up to 12.5 wt
%), making it a promising candidate as a hydrogen storage material.
Dating back to the period of 1985–1996, the groups of Saito,104−108 Cole-Hamilton,109 Shinoda,110 and Maitlis111 reported
on the production of hydrogen gas from methanol, where homogeneous
catalysts based on ruthenium, rhodium, or iridium were employed for
the dehydrogenation of anhydrous methanol either thermally or photochemically.
Depending on the catalyst and the reaction conditions, different products,
e.g., formaldehyde, formate salt, methylal (formaldehyde dimethyl
acetal), and methyl formate were obtained, influencing the yield of
hydrogen gas. Table 1 shows the thermodynamic parameters of different products formed
from the dehydrogenation of methanol. Dehydrogenation of methanol
to CO and H2 offers a possibility of high hydrogen storage
capacity (12.5 wt %), but the presence of CO is poisonous to fuel
cells (entry 1, Table 1). On the other hand, the formation of dimethoxymethane or methylal
offers a very low hydrogen storage capacity (2.1 wt %). For the purpose
of hydrogen storage, dehydrogenation of aqueous methanol, named "aqueous
methanol reforming", is attractive as it leads to the complete
dehydrogenation of methanol to inert CO2 (does not poison
the fuel cell catalyst), offering a high gravimetric hydrogen storage
capacity (12.0 wt %, entry 3, Table 1). More importantly, several examples for the direct
or indirect hydrogenation of CO2 to methanol have been
reported (vide infra, section 4) making methanol a material suitable for
the reversible storage of H2. A few reviews have been reported
for the production of H2 from methanol.112−114 Here, we focus on the aqueous methanol reforming processes catalyzed
by homogeneous transition metal complexes for the purpose of hydrogen
storage.</p><!><p>Major breakthroughs in the aqueous methanol reforming
process under mild catalytic conditions were reported in 2013 by the
groups of Beller115 and Grützmacher.77 Beller utilized ruthenium-MACHO pincer complexes
(Ru1, and Ru2) for the dehydrogenation of
MeOH/H2O mixture to CO2 (or CO32–) and H2 (Scheme 2).115 In the presence
of complex Ru1 or Ru2 and base, methanol
is first dehydrogenated to produce formaldehyde, which in the presence
of water is dehydrogenated to formic acid, and finally formic acid
is dehydrogenated to produce CO2. The overall process produces
3 equiv of hydrogen gas (CH3OH + H2O = 3H2 + CO2) at 65–95 °C, significantly
lower than the temperature of heterogeneously catalyzed aqueous methanol
reforming processes (>200 °C).125 Catalytic
conditions were optimized by varying the catalyst, concentration of
base, methanol–water ratio, and reaction temperature. Significant
catalytic activity was observed using the conditions of 8 M KOH, reaction
temperature of 91 °C, and a 9:1 MeOH/H2O solution,
where Ru1 (19 ppm) demonstrated a TOF of 1023 h–1, whereas Ru2 (19 ppm) demonstrated even a higher turnover
frequency (TOF) of 2668 h–1. Hydrogen gas was detected
in highly pure form; contaminants such as CO and CH4 were
detected in very small amounts (<1 ppm), lower than those reported
using heterogeneous catalysts.</p><p>A proposed mechanism
for the aqueous phase dehydrogenation of methanol
has been outlined in Scheme 3. The first step is the reaction of the precatalyst Ru2 with a base to produce the active species ruthenium amido
complex Ru2a. Complex Ru2a can dehydrogenate
methanol via an "outer-sphere" concerted process to
form HCHO and complex Ru2b via a transition state TSI, although HCHO was not detected in solution. Release of
H2 from Ru2b can regenerate complex Ru2a which can coordinate with HCHO. Attack of hydroxide on
the coordinated aldehyde forms a gem-diolate complex
which is presumably stabilized by solvent (MeOH/H2O) as
shown in the transition state TSII. Further elimination
of H2 from the gem-diolate complex leads
to the formation of a formate species, which can either decoordinate
to regenerate the active species Ru2a or release CO2 to produce complex Ru2b via the transition state TSIII. A third equivalent of H2 is released from Ru2b via metal–ligand cooperation, regenerating the
active species Ru2a. Notably, hydride species Ru2c, Ru2d, and Ru2e were observed in solution,
suggesting them to be the resting states. More detailed mechanisms
elaborating the role of the base, solvent, and metal–ligand
cooperation have been reported by several groups in the recent past.126−131</p><p>Around the same time, an anionic ruthenium complex Ru3 was reported by Grützmacher for the dehydrogenation
of CH3OH/H2O mixture to CO2/H2.
Using 0.5 mol % Ru3, 80% conversion of methanol was obtained
in 10 h at 90 °C (THF solvent) under neutral conditions (TOF,
54 h–1).77 Compared to
Beller's catalyst discussed above, Grützmacher's
catalyst exhibited a lower TOF but a higher methanol conversion. More
remarkably, catalysis was achieved under a neutral condition without
needing any base or additive. Furthermore, the practicality of the
system was demonstrated by feeding the evolved 3:1 H2/CO2 gas mixture directly to power an H2/O2 fuel cell. Based on the experimental studies, a mechanism involving
the noninnocence of the trop2dad ligand was proposed (Scheme 4). The catalysis
starts with the reaction of the anionic [Ru(H)(trop2dad)]
(Ru3) with H2O to form complex Ru3a with the elimination of 1 equiv of H2. This is followed
by O–H activation of CH3OH to form Ru3b and subsequent C–H activation of the coordinated methoxide
to form Ru3c and 1 equiv of aldehyde. The formed aldehyde
then reacts immediately with water to form methanediol. The amino
imine complex Ru3c can dehydrogenate methanediol to generate
formic acid, forming complex Ru3e via Ru3d and releasing HCOOH. Finally, the Ru(0) complex Ru3e reacts with a base and eliminates H2 to regenerate the
active species Ru3. HCOOH can finally be dehydrogenated
by complex Ru3a to produce CO2. More detailed
mechanisms have been reported by the groups of Yang,132 Hall,133 and de Bruin/Grützmacher134 using DFT calculations.</p><p>Around the same time, Beller also reported aqueous methanol reforming
under neutral conditions using a dual catalytic system—a ruthenium-MACHO
pincer complex Ru4 and a ruthenium bisphosphine dihydride
complex Ru5 (Scheme 5).116 The Ru-MACHO-BH complex
(Ru4) under a low catalytic loading (e.g., 5 Î¼mol)
was found to be inactive for the dehydrogenation of methanol/water
(methanol, 9.0 mL; water, 1.0 mL) in the absence of a base. However,
under a higher catalytic loading (95 Î¼mol of Ru4), a dehydrogenation reaction was observed, and the gas evolution
rate was found to reach 61 mL/h. The requirement of higher catalytic
loading was attributed to the enhanced rate of decomposition of formic
acid as the acid could poison the catalyst by strongly binding to
the vacant ruthenium site. To overcome this challenge, a dual-catalytic
system was used where the role of Ru-MACHO-BH catalyst Ru4 was to convert methanol to formic acid (step a in Scheme 2) and the decomposition of
formic acid was performed by Ru(H)2(dppe)2, Ru5 (Scheme 5). A significant catalytic activity (TOF3 h = 138
h–1) was obtained in the presence of catalysts Ru4 (5 μmol) and Ru5 (5 μmol), and
hydrogen gas at the rate of 60 mL/h was collected from a methanol
(9.0 mL)/water (1.0 mL). Notably, less than 8 ppm of CO gas was detected
by GC. Comparative experimental studies revealed that although both
catalysts Ru4 and Ru5 were capable of performing
both steps independently (Scheme 2)—conversion of methanol to HCOOH (step a) and
decomposition of HCOOH to CO2 and H2 (step b)—when
used together they exhibited a positive synergistic effect and resulted
in higher catalytic activity.</p><p>Soon after, Milstein reported
the ruthenium PNN catalyst Ru6 (Scheme 5) for the dehydrogenation of a methanol/water
solution.118 At 100 °C, using 0.025
mol % complex Ru6, about 80% yield of H2 gas
was obtained from
the basic (KOH, 2 equiv) solution of MeOH/H2O in toluene.
No dehydrogenation was observed under neat conditions, without adding
an external solvent such as THF or toluene which was proposed to solubilize
the catalytically active species. A quantitative amount of hydrogen
gas was produced at 115 °C in 24 h from formic acid (1 mL, 2.65
mmol) in the presence of 2 equiv of KOH using complex Ru6 (0.09 mol %) and KOtBu (1.1 mol %).
However, in the absence of a base (KOH), only a 25% yield of hydrogen
gas was obtained. Mechanistic studies revealed that the precatalyst Ru6 in the presence of a base forms the dearomatized complex Ru6a which can react with H2O to form a hydroxy
complex, Ru6c, or with HCOOH to form a formate complex, Ru6b (Scheme 5). H2O was found to be detrimental for the decomposition
of the HCOOH step, and therefore a higher catalytic activity in toluene,
in which a high concentration of the catalyst and very low water concentration
are expected, was observed. On the basis of this observation, an equilibrium
among complexes Ru6, Ru6a, and Ru6b as depicted in Scheme 5 was suggested. The catalyst was found to be remarkably reusable
and showed no noticeable loss in activity for nearly 1 month under
the conditions described above, and ∼1.53 g of methanol (MeOH
was added on regular intervals without isolation or purification of
the catalyst) was fully converted to H2 and CO2 by using 0.0024 g of catalyst Ru6 exhibiting a turnover
number (TON) of ∼29 000.</p><p>In 2016, Reek reported
aqueous methanol reforming using a ruthenium
complex (Ru7, Table 2, entry 10) under basic conditions to produce H2, formate, and carbonate.121 Complex Ru7 (12 Î¼mol) in the presence of 8 M KOH catalyzed the
dehydrogenation of a methanol:water solution (9:1 v/v) in dioxane
(total volume 30 mL), exhibiting a TOF of 55 h–1 at 82 °C for 4.5 h.</p><!><p>An iridium-based catalyst for
methanol dehydrogenation was reported by Crabtree in 2015.119 The iridium complex Ir1 (Scheme 6A) catalyzed the
dehydrogenation of dry methanol under basic conditions (6.7 M KOH)
exhibiting a TON of 3612 in 24 h at 91 °C and 81% yield of hydrogen.
Unlike previous systems, this process predominantly formed formate
rather than CO2 or carbonate (<5%). Remarkably, the
iridium catalyst Ir1 was stable in air and no additional
water was used in the catalysis. With regard to the aqueous methanol
reforming process, Fujita and Yamaguchi reported in 2015 the dehydrogenation
of aqueous methanol using iridium complexes Ir2, Ir2a, and Ir2b that could be reversibly interconverted
in water by changing the pH of the solution (Scheme 6B).79 Addition
of NaOH to the aqueous solution of Ir2a resulted in the
formation of complex Ir2 at pH 6.8. Further addition
of NaOH produced complex Ir2b at pH 12. The reverse process
was demonstrated by the addition of triflic acid (HOTf) to the solution
of Ir2b producing complex Ir2 at pH 6.8
and subsequently complex Ir2a at pH 2.7. The catalytic
aqueous methanol reforming reaction was found to work under basic
conditions suggestive of the requirement of a basic medium for the
generation of the active species. Indeed, in the presence of the complex Ir2a (0.5 mol %) and NaOH (0.5 mol %), 84% yield of hydrogen
gas was obtained from a methanol/water (1:4) solution under reflux
conditions for 20 h. A continuous supply of base was needed as the
generated CO2 dissolves in the reaction medium and lowers
the pH, transforming the active species Ir2b to the inactive
species Ir2a. As the need for a continuous supply of
a base to keep the catalyst active is a bottleneck for the commercialization
of this technology, Inagaki, Fujita, and co-workers have recently
reported steam reforming of methanol using vapor-phase flow technology.135 The modified anionic iridium bipyridonate (Ir-bpyd)
complex (Ir2b) was immobilized on a periodic mesoporous
organosilica to make a heterogeneous catalyst that could catalyze
the dehydrogenation of a methanol/water mixture in the vapor-phase
reaction without needing a base. The use of vapor-phase reaction methodology
was attributed to prevent CO2 to neutralize the anionic
iridium bipyridonate complex keeping the catalytic center active.
Detailed studies on the acid–base equilibrium of iridium complexes
containing dihydroxy-bipyridine ligands were reported by Himeda and
Fujita.136 A separate study by the same
group revealed that the position of the hydroxy group in the bipyridine
also plays a crucial role in the substrate activation process and
in the catalytic outcome.137</p><p>In a similar direction, Beller, in 2017 reported an iridium-PNP-MACHO
pincer complex Ir3 (Scheme 6A) for the catalytic dehydrogenation of aqueous
methanol under basic conditions.123 The
catalytic activity of the iridium pincer complex was also compared
with the analogous ruthenium complex Ru2 (Table 2, entry 2) and iron complex Fe1 (Table 2, entry 5) and was found to be significantly lower. The base concentration
was found to play a significant role in catalytic activity. For example,
using 0.5 M KOH and 4.18 Î¼mol of catalyst Ir3 resulted
in a TON of 1400 for the dehydrogenation of MeOH/H2O solution
(9:1, v/v) at 94 °C; however, in the presence of 8.0 M KOH a
TON of 1900 was achieved and the catalyst was stable for up to 60
h. In the presence of an excess of base, CO2 was trapped
as carbonate and only traces of CO2 were detected by GC.</p><p>Other than ruthenium and iridium, a couple of rhodium complexes
have also been utilized for the aqueous methanol reforming process.
Seminal work by Cole-Hamilton in 1987 revealed that aqueous methanol,
albeit in a lower yield, can be dehydrogenated using the [Rh(2,2′-bipyridyl)2]Cl complex Rh1 (Scheme 6A).109 Despite
the seminal report, rhodium complexes were not studied for the aqueous
reforming of methanol for a long time. Thirty years later, in 2017,
Zhou reported the catalytic activities of several rhodium complexes
for aqueous methanol reforming at 70 °C without adding any external
base, out of which [Cp*Rh(NH3)(H2O)2]2+Rh2 (Scheme 6A) was found to be the most active.124 A slight acidic pH of the solution improved
the catalytic activity, and the highest TOF of 83 h–1 was observed at pH 6 using a phosphate buffer.</p><!><p>Although precious metals have been at the forefront
of catalysis, their low abundance and high cost raise the concern
of sustainability and create the need for the development of base-metal
catalysts. Lately, there has been a substantial development in the
direction of homogeneous catalysts based on complexes of earth-abundant
metals for (de)hydrogenation reactions which have been reported in
several recent reviews.138−141 In the direction of aqueous methanol reforming,
the first homogeneous catalyst of an earth-abundant metal, the iron
pincer complex Fe1, was reported by Beller in 2013.117 However, the iron complex Fe1 (Scheme 7) was not as stable
as the analogous ruthenium complex Ru2 (Table 2, entry 2) discussed earlier,
and the catalytic activity was found to diminish over time. An outer-sphere
mechanism analogous to that discussed for the ruthenium pincer complex
(vide supra, Scheme 3) involving metal–ligand cooperation was proposed.</p><p>In 2015, Bernskoetter, Hazari, and Holthausen
reported a highly
active iron catalyst, Fe2 (Scheme 7), for aqueous methanol reforming under acidic
conditions.120 This is unlike most of the
homogeneous catalysts reported earlier that required a basic medium.
The Fe-PNP pincer catalyst Fe2 (Scheme 7) in the presence of a Lewis acid cocatalyst
(LiBF4) exhibited a TON up to 51 000 for the production
of H2/CO2 from a 4:1 methanol/water solution.
The TON reported here was the highest reported for either the base-metal
catalysts or base-free systems. In the absence of water, methanol
was dehydrogenatively coupled to form methyl formate. Screening of
several Lewis acids revealed that small or oxophilic cations (e.g.,
Li+, Na+) and weakly or noncoordinating anions
(e.g., PF6–, BF4–, and OTf–) were the most effective. The use of
Lewis acids in a catalytic amount allows the catalysis to occur under
milder conditions.</p><p>Another base-metal catalyst for aqueous methanol
reforming was
reported by Beller in 2017 using manganese complexes of PNP-type tridentate
ligands.122 A series of manganese complexes
were evaluated for the dehydrogenation of MeOH/H2O (9:1)
solution under basic condition (8 M KOH) at 92 °C, out of which Mn1, Mn2, and a combination of [Mn(CO)5Br] and 10 equiv of the ligand HN(CH2CH2)P(CH(CH3)2)2 (PNPiPr ligand) exhibited reasonable catalytic activities of TON5 h = 54, 65, and 68 respectively (Scheme 7). Notably, the catalysis in the case of Mn1 was found to be highly sensitive to light irradiation
and all the catalytic experiments were performed with the exclusion
of light. Catalyst Mn3 was also active but showed lower
catalytic activity (TON5 h = 41). Interestingly,
the presence of a Lewis acid (LiBF4) inhibited the catalysis,
unlike the reports by Bernskoetter, Hazari, and Holthausen.120 Addition of an excess of the PNPiPr ligand to complex Mn2 resulted in
remarkable stability for longer than a month, showing a TON of more
than 20 000. The stability of this system was higher than that
of the analogous Fe-PNP pincer catalyst Fe1 (Scheme 7) reported by Beller
for aqueous methanol reforming which was stable only for up to 5 days.
However, the catalytic activities of the Mn-PNP pincer catalysts were
still lower than that of the iron pincer catalyst Fe2 (Scheme 7) reported
by Bernskoetter, Hazari, and Holthausen, which exhibited a TON of
up to 51 000. A mechanism similar to that discussed earlier
in the case of Ru-PNP catalyst (Scheme 3) was proposed based on the NMR and ex situ IR investigations.</p><!><p>Similar to aqueous
methanol, aqueous formaldehyde also offers the
possibility of being used as a hydrogen storage material although
with a lower storage capacity. In a typical mechanism, formaldehyde
or paraformaldehyde can react with water to form methanediol, which
in the presence of a catalyst can liberate H2 to form HCOOH.
Further catalytic dehydrogenation of HCOOH can liberate H2 and CO2 (Scheme 8). Overall, the formaldehyde–water system offers a
hydrogen storage capacity of 8.4 wt %. Catalytic hydrogenation of
CO2 to formaldehyde has also been reported, making formaldehyde
a potentially reversible hydrogen storage material.142−146 However, unlike examples reported for the catalytic dehydrogenation
of methanol (section 2.1) and formic acid (section 2.4.4), dehydrogenation of aqueous formaldehyde
has received scant attention. In this section, we review homogeneous
catalysts that have been utilized for the dehydrogenation of aqueous
formaldehyde or paraformaldehyde (Table 3).</p><!><p>The structure of Ru9 is shown on the left.</p><!><p>Prechtl,
in 2014, published the first report on the aqueous phase
reforming of formaldehyde or paraformaldehyde in the presence of the
[(Ru(p-cymene))2(μ-Cl)2Cl2] complex Ru8 (Table 3, entry 1).147 Interestingly, the dehydrogenation process was also accomplished
in the absence of a base (although the presence of a base accelerates
the dehydrogenation process slightly), unlike in the cases of dehydrogenation
of methanol (vide supra, section 2.1) and formic acid (vide infra, section 2.4.4) where the presence of a base was found to be crucial for most of
the catalytic systems. The efficiency of the catalytic (0.1 mol % Ru8) dehydrogenation was found to be dependent on the reaction
temperature, and the best catalyst performance was achieved at 95
°C producing H2 in 84% yield in 60 min. The use of
a pH buffer during catalysis was beneficial but not crucial, and efficient
yields of H2 were obtained for a wide range of pH, e.g.,
pH 2.4 (75% H2) and pH 9 (73% H2) with a maximum
yield at pH 5.5 (85% H2). Catalyst recyclability and long-term
stability were also demonstrated, and continuous production of H2 gas was achieved by simply recharging the aqueous phase with
paraformaldehyde. The nature of the active catalyst was probed by
NMR spectroscopy, mass spectrometry, and isotope-labeling experiments
that suggested a dinuclear ruthenium species [(Ru(μ-cymene))2(μ-H)(μ-HCO2)μ-Cl]+ (Ru9, Scheme 8) to be the active catalytic species. Moreover, control experiments
using metal nanoparticles did not show any catalytic activity, suggestive
of homogeneous catalysis. Furthermore, in support of the nature of
the active catalyst, complex [(Ru(μ-cymene))2(μ-H)(μ-HCO2)μ-Cl]BF4 (Ru9) was independently
synthesized and tested in catalysis, resulting in a catalytic activity
similar to that of Ru8 (Table 3, entry 1). On the basis of the mechanistic
studies, the mechanism of the catalytic cycle was proposed to be analogous
to the one reported by Puddephatt for formic acid dehydrogenation
by a binuclear ruthenium complex.150</p><p>Deska and Prechtl utilized a similar ruthenium system for the dehydrogenation
of formaldehyde and its application for transfer hydrogenation reactions.151 Suenobu and Fukuzumi reported in 2015 the decomposition
of aqueous paraformaldehyde to H2 and CO2 using
an organoiridium complex Ir4 that is interconvertible
to Ir4a and Ir4b upon changing the pH as
shown in Scheme 9.148 At pH 11, a catalytic amount of complex Ir4b (5.0 mM) facilitated the decomposition of aqueous paraformaldehyde
(2.0 mg, 66.7 mmol) to produce H2 and CO2 (2:1
molar ratio) at 298 K with a TON of 21 (14 h). Reducing the catalyst
concentration to 1.0 mM did not change the TON, whereas a higher TON
of 51 was obtained at a higher temperature of 333 K. The rate of H2 production was found to decrease upon decreasing pH and no
H2 production was observed at pH 3, suggesting that the
hydroxo form Ir4 rather than Ir4a or Ir4b is the actual catalyst.</p><!><p>From ref (148). CC
by 3.0.</p><!><p>Similar to Ir4, Ir4a, and Ir4b complexes, Fujita and Yamaguchi
utilized the pH-dependent iridium
complexes Ir2, Ir2a, and Ir2b discussed earlier for aqueous methanol reforming (Scheme 6), for the dehydrogenation
of aqueous formaldehyde.79 Complexes Ir2, Ir2a, and Ir2b (0.5 mol %)
were tested for the dehydrogenation of formaldehyde–water solution
(37% formaldehyde) under reflux conditions for 20 h. Poor catalytic
activity was exhibited by complexes Ir2a and Ir2, but complex Ir2b showed higher catalytic activity
and 57% hydrogen gas was produced. Interestingly, the catalytic activity
was enhanced in the presence of an additional 0.5 mol % NaOH, resulting
in an 89% yield of hydrogen gas.</p><p>In the direction of producing
H2 gas from formaldehyde,
Trincado, de Bruin, and Grützmacher utilized the redox noninnocent
complex Ru3, used earlier for the dehydrogenation of
MeOH/H2O mixture (Scheme 4), and analogous complexes for producing H2 gas from a formaldehyde–water mixture (Scheme 10).78 Interestingly, the dehydrogenation product was trapped as a carbonate
salt, unlike the above-discussed examples where CO2 gas
was eliminated upon dehydrogenation of HCHO. Anionic complexes Ru3kc, Ru3ka, and Ru3kb (0.4 mol
%) that exist as a tight ion pair77 were
screened for the catalytic decomposition of aqueous formaldehyde (1
mmol, initial concentration = 0.47 M) at 60 °C in a water/THF
(10:1) mixture (Scheme 10). A low yield of hydrogen gas (32%) and a poor TON (115 in
12 h) were observed in the absence of a base using catalyst Ru3k. However, using an additional 6 equiv of KOH (keeping
the remaining conditions the same) resulted in the release of 86%
hydrogen gas and exhibited a high TON of 430 (in 2 h). The addition
of KOH was found to drive the reaction by trapping the evolved CO2 as K2CO3.</p><!><p>db18-C-6 = dibenzo-18-crown-6.</p><!><p>Although the redox-active complexes (Ru3ka, Ru3kb, Ru3kc, Scheme 10) exhibited high catalytic turnovers and
recyclability, the use of an excess base in catalysis limits its commercial
application due to the production of a stoichiometric salt waste in
the regeneration of formaldehyde. To avoid this problem, Ertem and
Himeda, in 2018, reported the ruthenium catalyzed dehydrogenation
of aqueous formaldehyde in the absence of a base. Remarkably, a record
TON (up to 24 000) was obtained without the use of any additive
or a base.149 Screening of several ruthenium
catalysts (10 μmol) at 50 °C for 7 h revealed that complex Ru10 (Scheme 11) is the most active catalyst for the dehydrogenation of aqueous
formaldehyde. The catalysis was found to be significantly dependent
on the temperature, and the optimum temperature was found to be 95
°C. With the use of 10 μmol of catalyst Ru10, 5 mmol of paraformaldehyde (in 5 mL of H2O) was dehydrogenated
at 95 °C to obtain a high yield and selectivity of H2 gas (95%) in 7 h.</p><!><p>Amine-boranes (RH2N·BH2R) are Lewis
acid and Lewis base adducts of amines and boranes, where protic and
hydridic hydrogen atoms are adjacent to each other, making the release
of hydrogen gas kinetically favorable. Low molecular weight amine-boranes
can offer very high gravimetric hydrogen capacities; for example,
H3N·BH3 can exhibit a hydrogen storage
capacity of up to 19.6 wt %. This has led to immense studies in pursuit
of efficient catalysts to dehydrogenate various amine-boranes as discussed
below.</p><!><p>Dehydrogenation of low molecular weight linear amine-boranes have
been intensively studied in the past decade as they exhibit high hydrogen
storage capacities and offer a potential route to make new types of
"B–N" polymer isosteres with polyolefins. Although
amine-boranes can be dehydrogenated thermally without needing a catalyst,
the use of a transition metal catalyst can allow the dehydrogenation
to occur under milder conditions. Moreover, a catalyst can control
the kinetics and influence the final product distribution, which is
crucial as it affects the hydrogen storage capacity. Table 4 summarizes the hydrogen storage
capacities of various linear amine-boranes and their dehydrogenation
products. As described in Table 4 and Scheme 12, H3B·NH3 can liberate 1, 2, or
>2 equiv of hydrogen gas, forming polyaminoborane [H2BNH2]n (6.48 wt %), borazine
[HBNH]3 (12.96 wt %), or polyborazylene (>12.96 wt %),
respectively.
A primary amine-borane such as H3B·NMeH2 can dehydrocouple to form N-methylated borazine [HBNMe]3 and liberate 2 equiv of H2 gas, offering a high storage
capacity of 8.86 wt %. Alternatively, H3B·NMeH2 can also dehydrocouple to form polyaminoborane [H2BNMeH]n, exhibiting a relatively lower
hydrogen storage capacity of less than 4.43 wt %. A secondary amine-borane
such as H3B·NMe2H can release only 1 equiv
of hydrogen gas forming the cyclic dimer [H2BNMe2]2 via the amino-borane H2B=NMe2 intermediate, offering a low hydrogen storage capacity of
3.38 wt % (Scheme 12). However, the products resulting from the dehydrocoupling of H3B·NMe2H are well-defined and soluble in common
organic solvents, unlike H3B·NH3 that upon
dehydrogenation usually results in a mixture of insoluble products.
This makes H3B·NMe2H a strategic choice
for mechanistic investigations via kinetics and spectroscopic studies.
Indeed, several pioneering studies aimed at elucidating the mechanisms
of the catalytic cycle for the dehydrogenation of amine-boranes have
been performed using H3B·NMe2H.152−158 Compared to N-substituted amine-boranes, reports on the dehydrogenation
of B-substituted amine-boranes are less common. Release of 2 equiv
of hydrogen gas from MeH2B·NH2R (R = H,
Me) can produce borazine derivatives [MeBNR]3 exhibiting
hydrogen storage capacities of 8.86 wt % (R = H) and 6.76 (R = Me)
wt %. Properties and catalytic dehydrogenation of these amine-boranes
have been reviewed earlier in detail and will not be discussed here.27−36</p><!><p>[M] = metal catalyst.</p><!><p>Despite the high gravimetric hydrogen
capacities of linear amine-boranes
and ease of dehydrogenation as reflected by several reports using
a range of catalysts, a practical application of amine-boranes for
hydrogen storage has not been demonstrated yet. This is mainly because
of a high thermodynamic barrier for the regeneration of charged fuel
(amine-borane) from the dehydrogenated product (spent fuel). Dixon
and Liu proposed that a hybrid system designed by combining both BN
and CC fragments might result in a material that can dehydrogenate
with a minimum thermodynamic overpotential as the dehydrogenation
of an alkane is endergonic whereas the dehydrogenation of amine-boranes
is highly exergonic.160 For example, the
dehydrogenation of H3B·NH3 is exergonic
by −13.6 kcal mol–1 while the dehydrogenation
of its alkane isostere H3C–CH3 is endergonic
by 23.9 kcal mol–1 (Scheme 13). Dixon and Liu also demonstrated that
a cyclic amine-borane such as 1,2-BN cyclohexane, containing both
BN and CC moieties, leads to a thermodynamically reversible hydrogen
storage pathway (taking into account the aromatic stabilization) with
ΔG(dehydrogenation) = 1.9 kcal mol–1 (Scheme 13).160</p><p>The first example of dehydrogenation
of cyclic amine-boranes was
reported by Liu in 2011.161 Liu reported
the synthesis of an air- and moisture-stable amine-borane, BN-methylcyclopentane,
and its dehydrogenation under thermal conditions to release 2 equiv
of hydrogen gas, forming the tricyclic borazine.161 Regeneration of the charged fuel BN-methylcyclopentane
from the spent fuel borazine was also demonstrated in an overall yield
of 92% using a two-step methodology (Scheme 14). Furthermore, BN-methylcyclopentane was
also demonstrated to be a potential candidate for hydrogen storage
by evaluating several relevant properties such as thermal stability
below 40 °C, a viscosity of 25 cP (comparable to ethylene glycol),
and a pure hydrogen stream produced upon thermal dehydrogenation.162</p><p>Furthermore, in 2011, Liu reported the synthesis
and dehydrogenation
of 1,2-BN cyclohexane and its thermal dehydrogenation at 150 °C
to produce tricyclic borazine (Scheme 15).163 Later,
in 2016, Weller and Liu utilized iridium and rhodium bis-phosphine
complexes to perform dehydrogenation under ambient conditions.164 Both BN-methylcyclopentane and 1,2-BN cyclohexane
were demonstrated to exhibit a hydrogen storage capacity of 4.7 wt
%, although further dehydrogenation of the cyclic backbone can result
in a higher hydrogen storage capacity of up to 9.4 wt %. In a similar
direction, Liu reported the synthesis of bis-BN cyclohexane and its
dehydrogenation to form a cagelike structure catalyzed by the ruthenium-PN
complex Ru11 (0.5 mol %) at 65 °C (Scheme 15).165</p><!><p>Difficulty in the regeneration of amine-boranes (charged
fuel) from the spent "BN" fuel poses a significant challenge
in the practical utility of amine-boranes. Multistep processes involving
stoichiometric reagents (generating waste) for the regeneration of
H3B·NH3 from the waste BNHx products have been demonstrated.166−169 Gordon has reviewed regeneration of H3B·NH3 using a multistep protocol consisting of (a) digestion (H+ addition), where a spent fuel (BNHx)
is protonated by a strong acid or a weak acid, e.g., alcohol, amine,
or thiol (HX) to produce BX3 and NH3; (b) reduction
(H– addition), where BX3 reacts with
a reductant, e.g., metal hydride, in the presence of an amine to form
H3B·NR3; and (c) ammoniation, where the
reaction of H3B·NR3 with NH3 regenerates H3B·NH3.170 Remarkably, Sutton and Gordon reported in 2011 a one-pot
regeneration of H3B·NH3 from the spent
fuel polyborazylene by the reaction of anhydrous N2H4 in liquid NH3.171 Polyborazylene
(BNHx) reacted with 1.35 equiv of anhydrous
N2H4 at 40 °C in a sealed vessel for 24
h to regenerate H3B·NH3 (92% yield) and
N2 as the byproduct. However, the use of N2H4, which is a potential hydrogen storage material in itself,
for the regeneration process presented a new set of practical challenges
in terms of handling hydrazine such as toxicity, instability, and
risk of explosion as well as high cost and multistep synthesis of
N2H4 from NH3. Interestingly, Manners
reported that the B=N bond of a spent fuel, e.g., H2B=NiPr2, can be converted
back to the charged fuel H3B·NiPr3 via transfer hydrogenation using amine-boranes,
e.g., H3B·NHRR′ (R, R′ = H or Me) or
linear diborazane Me3N·BH2·NHMe·BH3.172,173 Remarkably, the reaction occurred
without needing any catalyst at 20 °C (THF solvent) to give yields
up to 90%. However, this concept too does not offer a practical sustainable
solution as the regeneration process produces spent BN fuel and requires
sacrificial charged fuel. Furthermore, Manners reported that aminoboranes
H2B=NR2 can be converted to H3B·NHR3 in the presence of H2O in a single
step without needing a catalyst. The reaction was driven by the formation
of insoluble borate products (BxOyHz). However,
only a low yield of ∼30% was achieved. A higher yield was achieved
by the addition of sacrificial agents, e.g., BH3·THF
or LiBH4. Thus, it is essential to develop an efficient
and atom-economical route to regenerate amine-boranes from the spent
BN fuel. There is no report on the direct hydrogenation of BN spent
fuel such as borazine to regenerate H3B·NR3. As H3B·NR3 can be easily dehydrogenated
at high temperatures, the use of low temperature for the hydrogenation
reaction is recommended.169 Timoshkin reported
that the use of a Lewis acid can significantly decrease the activation
energy of hydrogenation.174 Szymczak and
Heiden have performed DFT calculations and suggested that coordination
of borazine to a transition metal fragment such as M(CO)3 can decrease the activation energy for the reactivity of borazine
toward a hydride source.175</p><!><p>To overcome
the challenges of regenerability and reversibility,
liquid organic hydrogen carriers (LOHCs) have been proposed as potential
hydrogen materials. LOHCs are low molecular weight organic compounds
in the liquid state at room temperature that can liberate hydrogen
gas (fuel) in the presence of a catalyst, forming spent fuel that
can be converted back to the charged fuel by hydrogenation, thus closing
the loop.176 The liquid state gives the
advantage of utilizing the established infrastructure for delivering
gasoline fuels. A model for the utilization of LOHC for the hydrogen
economy is shown in Figure 1. As depicted, an LOHC ideally of low flammability and toxicity
and in the liquid state, can be easily transported to a fuel station
where it can be stored indefinitely under ambient conditions. H2 can be produced at the fuel station via catalytic dehydrogenation
of the LOHC and can be used for energy generation in a fuel cell or
for industrial applications. For transportation purposes, LOHCs can
be directly loaded to a hydrogen vehicle where onboard catalytic dehydrogenation
can produce H2 gas that can be fed to a fuel cell to drive
the vehicle. The spent fuel (preferably liquid) produced in the dehydrogenation
process can then be transported conveniently to a suitable hydrogenation
facility where H2 produced from renewable sources can be
used to hydrogenate the spent fuel back to the LOHC (charged fuel
carrier). Reviews on various perspectives of LOHCs, e.g., chemical
and economic properties or supply chain strategies, have been reported
earlier.176−182 Here, we focus on the recent developments in LOHCs facilitated by
homogeneous transition metal catalysts.</p><!><p>Model depicting use of
LOHCs in hydrogen economy.</p><!><p>Carbocycles
contain several interesting properties such as high gravimetric storage
capacities, being liquid at room temperature, high boiling points,
low toxicity, CO-free dehydrogenation products, and high abundances,
making them suitable candidates for LOHCs. The major drawback associated
with carbocycles is the high barrier for the dehydrogenation reaction,
requiring harsh reaction conditions for the dehydrogenation process
such as a high temperature, 200–350 °C. This makes it
challenging for using homogeneous catalysts for the dehydrogenation
process, and most of the demonstrations of carbocycles as LOHCs involve
heterogeneous catalysts. Early studies using heterogeneous catalysts
were focused on LOHCs based on methylcyclohexane/toluene (6.1 wt %
theoretical hydrogen storage capacity),183,184 cyclohexane/benzene (7.14 wt %),185 decalin/naphthalene
(7.3 wt %),186 and perhydro-dibenzyltoluene/dibenzyltoluene
(6.2 wt %).187 Remarkably, LOHC technologies
based on methylcyclohexane/toluene and perhydro-dibenzyltoluene/dibenzyltoluene
were recently commercialized by Chiyoda Corp.188 and Hydrogenious LOHC Technologies,189 respectively. However, issues of selectivity during (de)hydrogenation
reactions due to the possibility of the formation of multiple intermediates
remain to be sorted out, offering scope for the discovery of new and
more efficient LOHCs.</p><!><p>Nitrogen-
or oxygen-containing heterocycles are thermodynamically advantageous
for dehydrogenation reactions compared to the corresponding carbocyclic
compounds.177,190,191 For practical advantages heterogeneous catalysts have been investigated
extensively to develop LOHCs based on heterocycles. A seminal discovery
in this direction was made by Pez and co-workers, who developed an
LOHC based on N-ethylcarbazole (NEC) with a hydrogen
storage capacity of 5.8 wt %.192 Although
a facile dehydrogenation and hydrogenation process has been demonstrated,
the NEC/H12–NEC system has some drawbacks that limit
its commercialization such as the solid state of N-ethylcarbazole and the formation of unwanted side products during
dehydrogenation. Jessop and co-workers193 and Ke and Cheng194 independently reported
an LOHC system based on octahydro-1-methylindole/1-methylindole constituting
a hydrogen storage capacity of 5.8 wt %. In a similar direction, LOHC
based on 2-[(N-methylcyclohexyl)methyl]piperidine/2-(N-methylbenzyl)pyridine with a hydrogen storage capacity
of 6.1 wt % was reported by Suh and Park.195 Kempe discovered an LOHC based on phenazine (PHZ) that can be synthesized
from cyclohexane-1,2-diol obtained from hydrogenolysis of lignin.196 With the use of a bimetallic catalyst, Pd2Ru@SiCN, PHZ was hydrogenated to form 14H-phenazine, making
it a hydrogen storage material of 7.2 wt %. Milstein recently reported
a solvent-free LOHC based on 2-picoline/2-methylpiperidine (6.1 wt
%) and 2,6-lutidine/2,6-dimethylpiperidine (5.3 wt %) using a palladium-based
heterogeneous catalyst generated in situ, efficient
for both hydrogenation and dehydrogenation reactions under relatively
mild conditions.197 In addition to the
demonstration of LOHCs by heterogeneous catalysts, there are a couple
of examples in which molecular complexes were employed. Fujita and
co-workers reported the use of an iridium complex for the dehydrogenation
of 2,5-dimethylpiperazine to 2,5-dimethylpyrazine releasing 3 equiv
of H2 in the presence of a very small amount (0.5–1
mL) of solvent (Scheme 16).198 Under the catalytic conditions
of Ir2 (Scheme 16, 0.25 mol %) and 6,6′-dihydroxy-2,2′-bipyridine
(L1, 0.50 mol %) in the p-xylene/water
solvent, quantitative dehydrogenation of dimethylpiperazine to 2,5-dimethylpyrazine
was observed. Moreover, hydrogenation of 2,5-dimethylpyrazine to 2,5-dimethylpiperazine
took place quantitatively (Scheme 16). Furthermore, closed-loop conversion between 2,5-dimethylpyrazine
and 2,5-dimethylpiperazine was also demonstrated with the quantitative
yields at least four times.</p><!><p>A new approach in the area of
LOHCs was revealed by Milstein and
co-workers by utilizing the concept of acceptorless dehydrogenative
coupling of amines and alcohols to form amides199,200 with the liberation of H2 and its reverse reaction—hydrogenation
of amides to form alcohols and amines.200 In 2015, Milstein reported the first LOHC based on this concept
involving the dehydrogenative self-coupling of the widely used, industrial,
2-aminoethanol (AE) and the reverse hydrogenation reaction exhibiting
a theoretical hydrogen storage capacity of 6.56 wt % (Scheme 17).201 Using a combination of the ruthenium pincer complex Ru12 (0.5 mol %) and KOtBu (1.2 mol %), AE
(85% conversion) was dehydrogenated to form the cyclic glycine anhydride
(GA, 60%) as the major product and linear peptides (LP) as the side
product. H2 gas was generated in 77% yield. The same catalyst
system catalyzed the hydrogenation of glycine anhydride under 50 bar
H2 at 110 °C, 48 h, in approximately quantitative
yield. Importantly, the mixed products GA + LP produced from the dehydrogenation
reaction were also successfully hydrogenated under the same catalytic
conditions, indicating that the formation of linear peptides is not
a problem for the regeneration of AE. Complex Ru13 also
catalyzed the same dehydrogenation and hydrogenation reactions although
with lower yield and selectivity.</p><p>Furthermore, Milstein
reported in 2016 an LOHC based on the dehydrogenative
coupling of ethylenediamine and ethanol to form N,N′-diacetylethylenediamine (DAE) using the
ruthenium pincer catalyst Ru6 (Scheme 18).202 The same
complex was also found to catalyze the reverse reaction, i.e., the
hydrogenation of DAE to a mixture of ethylenediamine and ethanol,
forming an LOHC system with a theoretical hydrogen storage capacity
of 5.3 wt %. Under the catalytic conditions of complex Ru6 (0.02 mmol) and KOtBu (0.024 mmol),
5 mmol of ethylenediamine and 12 mmol of ethanol were dehydrogenated
(100% conversion) in refluxing dioxane (2 mL) to form DAE in 93% yield
and N-(2-aminoethyl)-acetamide (AEA) in 7% yield.
Similarly, in the presence of complex Ru6 (1 mol %) and
KOtBu (1.2 mol %), DAE was hydrogenated
(50 bar H2, 115 °C) to form a mixture of ethylenediamine
(91%), AEA (9%), and ethanol (70% yield). Repetitive demonstration
of dehydrogenation/hydrogenation cycles was also demonstrated without
adding more catalyst.</p><p>Another LOHC based on alcohol
and amine, with a theoretical hydrogen
storage capacity of 6.66 wt %, was reported by Milstein in 2018 using
1,4-butanediol and ethylenediamine (Scheme 19).203 Employing
the ruthenium pincer complex RuPPh2NNH (Ru14, 1 mol %) in the presence of 2 mol % KOtBu (2 mL 1,4-dioxane, 120 °C), catalytic dehydrogenation
of 1,4-butanediol (1 mmol) and ethylenediamine (0.6 mmol) resulted
in a mixture of bis-cyclic imide (70%), lactone (12%), and oligoamides.
Pure H2 gas (82 mL), as confirmed by GC, was also collected.
An analogous complex RuPtBu2NNH (Ru12, 1 mol %) in the presence of KOtBu (3 mol %) catalyzed the hydrogenation
(40 bar H2, 135 °C, 40 h) of the bis-cyclic imide
to regenerate 1,4-butanediol and ethylenediamine in approximately
90% yields.</p><p>Along a similar direction,
Prakash reported an LOHC based on "amine
reforming of methanol" in analogy with the steam reforming
of CH3OH using a combination of N,N′-dimethylethylenediamine (DMEDA) and CH3OH that can exhibit a hydrogen storage capacity of up to 5.3 wt %
(Scheme 20).204 With employing the ruthenium complex Ru4 (1 mol %) and K3PO4 (4 mol %), the
coupling of DMEDA (1 mmol) and methanol (4 mmol) at 120 °C for
24 h in toluene resulted in the formation of the corresponding diamide
(75%) and monoamide (22%) along with hydrogen gas (86%). The undesired
CO gas was detected by GC in 2.8% yield, suggested to originate from
decarbonylation of intermediate formaldehyde, in competition with
amine attack to form the α-amino alcohol intermediate. Interestingly,
using the RuPNPiPr complex Ru2 (1 mol %) and K3PO4 (5 mol %), keeping
the remaining conditions the same, 90% yield of hydrogen gas was observed,
and no CO was detected in the gas mixture. In order to study the reversibility
of the reaction, the crude reaction mixture after dehydrogenation
using catalyst Ru2 was subjected to hydrogenation conditions
(60 bar H2, 120 °C, 24 h) without adding any additional
catalyst. After 24 h, the formation of DMEDA was observed in 95% yield.
Furthermore, both dehydrogenation and hydrogenation reactions were
performed under neat conditions without adding any solvent and gave
good to moderate yields (76 and 60%, respectively). Recycling of the
catalyst up to three times was also demonstrated.</p><p>In the direction of using earth-abundant metals
in catalysis, Liu
and co-workers have recently demonstrated the proof of concept of
the same LOHC (DMEDA/CH3OH and diamide) using manganese-based
catalysts (Scheme 20).205 The best dehydrogenation activity
was achieved using the complex Mn4, whereas the complex Mn5 exhibited the highest yield for the hydrogenation reaction.
Compared to the ruthenium catalyst Ru2, a higher loading
(2 mol %) and higher temperature (165 °C for 16 h) were used
in the case of Mn4. Furthermore, a higher amount (∼5%)
of CO gas was detected in the case of Mn4 which could
be suppressed by the addition of Mn4 in two equal portions.</p><p>Based on control experiments and earlier reports, a plausible mechanism
was proposed for the manganese- or ruthenium-catalyzed dehydrogenative
coupling of diamine and methanol (Scheme 21). Methanol first is dehydrogenated to formaldehyde
that reacts with an amine moiety of DMEDA to generate the hemiaminal
species A1 that eliminates one molecule of H2 to afford the monoamide intermediate A2. A2 reacts with another molecule of formaldehyde to produce the final
diamide product A3. Another possible pathway could be
via the formation of methyl formate (A4) from methanol
that undergoes aminolysis with DMEDA to form the diamide A3. Fast condensation of formaldehyde with the amino groups and monoamide A2 circumvents the decarbonylation of formaldehyde to CO and
H2.</p><!><p>[M] = Ru or Mn catalysts.</p><!><p>Besides using a combination of alcohols and amines,
simple alcohols
have also been demonstrated as LOHCs. Recently, Fujita has reported
a reversible hydrogen storage system based on 1,4-butanediol/γ-butyrolactone206 catalyzed by the iridium complex Ir2 reported earlier for the development of the LOHC based on 2,5-dimethylpiperazine/2,5-dimethylpyrazine
(Scheme 16).198 Under the catalytic condition of 0.1 mol % Ir2, a solution of 1,4-butanediol in 1,2-dimethoxyethane under
reflux for 20 h resulted in dehydrogenation to produce γ-butyrolactone
and H2 gas in quantitative yields (Scheme 22). Solvent screening showed that anisole
was a better solvent in which a complete dehydrogenation was obtained
in only 3 h. Interestingly, the reaction also proceeded under neat
conditions, and almost complete conversion of 1,4-butanediol (15 mmol)
to γ-butyrolactone was obtained in the presence of 0.5 mol %
catalyst Ir2 and hydrogen gas was collected in 99% yield.
Moreover, the reverse reaction, i.e., hydrogenation of γ-butyrolactone
to 1,4-butanediol, was also demonstrated under neat conditions using Ir2 and ligands such as 6,6′-dihydroxy-2,2′-bipyridine
(L1) and triethylamine. Furthermore, successive interconversions
between γ-butyrolactone and 1,4-butanediol were also demonstrated
under neat conditions, and both the dehydrogenation and hydrogenation
steps were accomplished with almost quantitative yields. However,
the theoretical hydrogen storage capacity of this system is lower
than those of the systems reported using alcohols/amines and is limited
to 4.4 wt %.</p><!><p>L1 = 6,6′-dihydroxy-2,2′-bipyridine.</p><!><p>In a recent advance, Milstein and co-workers
reported an LOHC system
based on ethylene glycol (EG).207 EG is
industrially produced on a large scale with a global annual production
of 34 million tons, and it has applications in the automobile and
polyester industries.208 Additionally,
EG can be produced from biomass-derived hydrocarbons, making it a
sustainable and renewable LOHC.209 However,
dehydrogenation of ethylene glycol presents additional challenges
such as catalyst poisoning by chelation of the vicinal 1,2-diol; the
formation of α-keto ester byproduct, which can readily decompose
to aldehyde and CO which could again poison the catalyst; and formation
of unwanted cyclic side products such as 1,3-dioxolan-2-ylmethanol,
resulting in lower hydrogen capacities. However, remarkably, with
the use of the acridine-based ruthenium PNP complex Ru15,210,211 these difficulties were circumvented to
some extent. In the presence of Ru15 (1 mol %) and KOtBu (1 mol %), ethylene glycol (2.0 mmol)
in toluene/DME (1.0 mL/1.0 mL) at 150 °C and 72 h was dehydrogenated
to produce 54 mL of hydrogen gas, 2-hydroxyethyl glycolate (HEG) (33%),
and oligoesters (EG conversion 94%) (Scheme 23). Interestingly, dehydrogenation of EG
proceeded smoothly in the presence of the dearomatized ruthenium PNP
complex Ru16 (1 mol %) under base-free conditions (keeping
the remaining conditions the same), producing 61 mL of hydrogen gas
and oligoesters of up to hexamer, with the conversion of EG being
97%. Interestingly, the reaction mixture obtained from the dehydrogenation
process was hydrogenated back to ethylene glycol in excellent yield
(92%) by using Ru16 (1 mol %) and 40 bar H2 (48 h) in toluene/DME solvent, demonstrating the utility of EG as
an LOHC. Moreover, the reaction was also demonstrated under solvent-free conditions at a reduced pressure (95 mbar,
EG, 35.9 mmol, 2 mL) at 150 °C that resulted in 94% conversion
of EG in 7 days.</p><p>A catalytic cycle was proposed
based on the DFT calculations (Scheme 24). As outlined
in Scheme 24, the
first step is the coordination of EG to complex Ru16 to
form intermediate Ru16a, followed by
the dehydrogenation step via Ru16TSAB (24.7 kcal·mol–1) resulting in the
formation of a κ2-alkoxide coordinated Ru16b (−5.1 kcal·mol–1). Hemilability of
the hydroxy group facilitates the β-hydride elimination step
via Ru16TSBC (7.2 kcal·mol–1) forming a coordinated hydroxyaldehyde intermediate Ru16c (5.9 kcal·mol–1). Attack of another
molecule of EG on the coordinated aldehyde followed by dehydrogenation
via a concerted Zimmerman–Traxler-type six-membered transition
state (Ru16TSCD, 23.4 kcal·mol–1) results in the formation of intermediate Ru16d (3.7 kcal·mol–1). A similar decoordination
of the hydroxy group and β-hydride elimination step results
in the formation of metal-bound ester Ru16e (6.8 kcal·mol–1) which can regenerate complex Ru16 by
dissociation of 2-hydroxyethyl glycolate (HEG). Subsequently, HEG
can generate an oligoester through similar cycles.</p><!><p>All values correspond
to
Gibbs free energies at 423.15 K (in kcal·mol–1 with respect to the starting material). Reproduced with permission
from ref (207). Copyright
2019 Springer Nature.</p><!><p>Formic acid (FA)
has attracted significant interest as
a potential hydrogen storage material due to several qualities. For
example, FA is a kinetically stable liquid at room temperature and
is produced on a large scale by the chemical industry as well as by
biomass fermentation. Although the theoretical hydrogen storage capacity
(4.4 wt %) is lower than the target set by the U.S. Department of
Energy for 2020 (5.5 wt %), the dehydrogenation process to produce
CO2 is thermodynamically favorable (ΔG = −32.9 kJ/mol) at room temperature. Several examples have
been reported for the catalytic dehydrogenation of HCOOH, and this
topic has also been reviewed multiple times in the past.114,212−221 Laurenczy and Beller, in 2018, published a detailed review on the
dehydrogenation of formic acid catalyzed by both noble metals and
non-noble metals.114 Zell and Langer, in
2019, reviewed the dehydrogenation of formic acid catalyzed by both
homogeneous and heterogeneous catalysts.218 Huang published an update on the dehydrogenation of formic acid
using homogeneous catalysts in 2020.220 Two review articles, one by Yao and co-workers222 and another by Himeda,223 have
been published in 2021 on the catalytic dehydrogenation of formic
acid. The reverse reaction—direct hydrogenation of CO2 to HCOOH—has also been well-reviewed in the past.224,225</p><p>Different from dehydrogenation, hydrogenation of CO2 to HCOOH is thermodynamically uphill. Three approaches have been
used to overcome the thermodynamic barrier for the hydrogenation of
CO2 to HCOOH: (a) use of a stoichiometric base to stabilize
HCOOH as a formate salt, (b) use of a polar solvent such as DMSO to
stabilize HCOOH by H-bonding,226,227 and (c) use of ionic
liquids containing basic anions such as 1,3-propyl-2-methylimidazolium
formate to improve the solubility of CO2 and stabilize
the product HCOOH.228 However, it is important
to note that for a reversible process, as required for an LOHC, the
additive used or the byproduct formed during hydrogenation should
be compatible with the dehydrogenation reaction or vice versa in order to avoid purification and generation of waste. For example,
dehydrogenation of formic acid can be carried out in an acidic medium,
but CO2 hydrogenation might require a basic medium. Thus,
the development of compatible catalysts and the choice of an appropriate
hydrogen storage couple becomes highly important. Here we present
a synopsis of the reports demonstrating HCOOH as a reversible hydrogen
storage system catalyzed by a transition metal complex.</p><p>As the
regeneration of HCOOH from CO2 poses a thermodynamic
challenge, Papp and Joó reported a reversible hydrogen storage
system based on aqueous HCOONa that was dehydrogenated to NaHCO3 in the presence of a water-soluble catalyst [RuCl(mtppms)2]2 (mtppms = sodium diphenylphosphinobenzene-3-sulfonate),
originally developed by Laurenczy and co-workers.229 No additive or tuning of pH was required for this reversible
process. Hydrogenation reaction was achieved at 100 bar H2, at 80 °C, and dehydrogenation or decomposition reaction was
performed in a glass pressure tube at 80 °C. Almost no CO gas
(≤10 ppm) was detected in the gas phase by GC.</p><p>In the
same direction, Beller and co-workers reported a reversible
hydrogen storage system based on dehydrogenation of aqueous formate
solution to a bicarbonate salt using a ruthenium/phosphine catalyst
and the reverse reaction, i.e., hydrogenation of bicarbonates to formates.230 A catalytic combination of [RuCl2(benzene)]2/dppm (1:4 ratio) (dppm = diphenylphosphinomethane)
was used for the hydrogenation of several alkali metal bicarbonates
out of which NaHCO3 gave the best result where HCOONa was
formed in 96% yield (TON = 1108) in 20 h at 70 °C under 80 bar
H2 pressure (Scheme 25). The addition of external CO2 pressure
enhanced catalytic activity. Moreover, dehydrogenation of various
formates was explored using Ru/dppm at 60 °C in a DMF/H2O solution. Production of hydrogen gas was observed in all cases
without detection of any CO gas. The use of excess water assisted
in the liberation of H2 from formates that shifts the pH
of the solution to a more basic medium. The basic solution could trap
the generated CO2 and allow it to precipitate as bicarbonate
during the reaction. Under the catalytic conditions of 5.0 mmol of
[{RuCl2(benzene)}2] (Ru17) and
30 mmol of dppm (Ru/P = 1:6), 20 mmol of HCOONa was dehydrogenated
at 40 °C to produce 490 mL of hydrogen gas with a TON (3 h) of
2000. A limitation of the above-described process reported by Beller
is the use of different solvent systems for dehydrogenation (THF/H2O) and hydrogenation (DMF/H2O) reactions, which
poses a practical challenge in the utilization of this system for
reversible hydrogen storage. Moreover, the theoretical hydrogen storage
capacity of the HCOONa and H2O/NaHCO3 system
is only 2.32 wt %.</p><p>In 2009, Nozaki reported the hydrogenation of CO2 to
formate in a basic medium catalyzed by the iridium trihydride pincer
complex [Ir(H)3(PiPrNPyPiPr)] (Ir5, Figure 2).231 The catalytic activity of this complex
was the highest of its time (TOF 150 000 h–1 at 200 °C, 50 bar, and TON 3 500 000 at 120 °C).
The same complex was also later utilized for the dehydrogenation of
formic acid, albeit with a relatively lower catalytic activity (TON43 h 890 at 60 °C) compared to that of the hydrogenation
of CO2.232 Addition of base
significantly enhanced the catalytic activity, e.g., TON5 h of 4900 at 60 °C in the case of triethanolamine as a base.
Although the reversibility of the HCOOH/CO2 + H2 system was studied, the demonstration of a charging/discharging
cycle for the purpose of using formic acid as an LOHC was not performed
in this study. Following this discovery, several homogeneous catalysts
were reported for the demonstration of HCOOH/CO2 as a reversible
hydrogen storage system.</p><!><p>Catalytic conditions for the reversible conversion
of CO2 to HCOOH.</p><!><p>In 2012, Laurenczy and Beller reported a reversible conversion
of CO2 to formic acid (FA) catalyzed by a ruthenium complex
[RuCl2(benzene)]2 (Ru17) and bisphosphine
(diphenylphosphinoethane (dppe) or diphenylphosphinomethane (dppm))
ligand.233 Hydrogenation of CO2 to formic acid was accomplished using [RuCl2(benzene)]2 (10 mmol) and dppm (6 equiv) under a basic solution (NEt3) and a moderate pressure (30 bar H2 + 30 bar CO2). A reaction temperature of 100 °C was required for
the generation of the active species, after which the catalysis could
proceed at room temperature. This strategy was also applied to demonstrate
hydrogen loading and unloading to establish a "hydrogen battery".</p><p>Soon after, Hull, Himeda, and Fujita reported a reversible hydrogen
storage system catalyzed by an iridium complex using CO2, formate, and formic acid under mild conditions.80 Drawing inspiration from hydrogen bonding in nature and
the use of bases to relay protons in enzymatic catalysis, the authors
developed new iridium complexes based on proton-responsive ligands
bearing pendent-base moieties that could assist the interaction of
H2, CO2, H2O, HCO2H, HCO2–, and H+ in the primary coordination
sphere of iridium. Complex Ir6, synthesized using the
bridging ligand 4,4′,6,6′-tetrahydroxy-2,2′-bipyrimidine,
rapidly hydrolyzes in water to produce Ir6a (Scheme 26). Increasing the
pH to more than 5 results in deprotonation of the ligand phenol moieties
and the formation of Ir6b. Both catalysts Ir6a and Ir6b were active for storing and releasing H2 under mild conditions. In the presence of Ir6b as the catalyst, a 1:1 H2/CO2 mixture (0.1
MPa) was converted to 0.36 M formate (TON 7200) at 25 °C (pH
8.4) with an initial TOF of 64 h–1. This activity
was found to be superior to those of Ir7 (initial TOF
7 h–1) and [{Ir(Cp*)(OH2)}2(bpym)]Cl2 (initial TOF 0 h–1), suggesting
that the pendent hydroxy base plays an important role in the hydrogenation
process. The reverse reaction, i.e., the dehydrogenation of formic
acid to CO2, was studied under acidic conditions. Under
pH 3.5, complex Ir6b is partially protonated and exists
in an intermediary form between Ir6b and Ir6a. Dehydrogenation of an aqueous solution of HCOOH/HCOONa at pH 3.5
by partially protonated Ir6b produces a 1:1 mixture of
H2/CO2 with a remarkably high TOF (228 000
h–1 at 90 °C and TON of 308 000 at 80
°C).</p><p>Along the direction of pH-dependent
charging and discharging of
formic acid, Fukuzumi, in 2013, reported the hydrogenation of CO2 using an iridium complex (Ir4a) in a weakly
basic medium (pH 7.5 in H2O) at ambient temperature and
pressure.234 The dehydrogenation of formic
acid was demonstrated at pH 2.8 (acidic medium) by using Ir4a at room temperature (Scheme 27).</p><p>Pidko in 2014 reported the reversible hydrogenation of CO2 using the ruthenium PNP pincer catalyst Ru18 (Figure 2).235 Optimization of catalytic conditions for the dehydrogenation
of FA revealed that a catalytic combination of Ru18 and
NR3 (R = Et, hexyl) in DMF solvent is highly efficient.
A catalytic combination of Ru18 (1.42 Î¼mol) and
NEt3 (33.5 mmol) in DMF (35 mL total volume) exhibited
significant dehydrogenation of FA with TOF of 257 000 h–1 at 90 °C (TON = 326 500). Catalytic hydrogenation
of CO2 to formate under a basic medium was also demonstrated
by using the complex Ru18 in DMF. The combination of Ru18 and 1,8-diazabicyclo[5.4.0]undec-7-ene (DBU) at 120 °C
(H2, 30 bar, and CO2, 10 bar) resulted in a
TOF of 1 100 000 h–1, which is the
highest reported catalytic activity for the hydrogenation of CO2 to date, higher than the earlier record value reported by
Nozaki231 using the Ir pincer complex Ir5 and that by Fujita using Ir6b (Figure 2).80 The nature of the base was found to significantly affect the catalytic
outcome. For example, (a) a higher TON was obtained (TOF = 256 000
h–1 and TON = 706 500) in the case of a less
volatile amine such as trihexylamine, in comparison to that of NEt3 for FA dehydrogenation; (b) a stronger base resulted in a
higher AAR (acid to amine ratio) at higher temperature and shorter
reaction time; and (c) the rate-determining step of the catalytic
cycle was found to be influenced by the strength of the base (e.g.,
DBU and NEt3).</p><p>In a similar direction, Plietker and
co-workers in 2014 reported
a reversible hydrogen storage system based on amine/CO2 using a Ru-PNNP catalyst (Ru19) (Scheme 28).236 The charged system could be stored for several days without loss
of efficiency, and several hydrogenation and dehydrogenation cycles
were demonstrated without changing the catalyst or the reaction vessel.
Efficient hydrogenation of CO2 (dry ice) was accomplished
by using catalyst Ru19. Using 0.015 mol % Ru19 and 65.7 mmol of DBU, at 100 °C and 70 bar H2, 20
g of dry ice was hydrogenated to the corresponding DBU formate salt
in 4 h (yield 84% and TON = 5600). Interestingly, at a higher catalyst
loading (0.075 mol % Ru19), a bis-formic acid DBU adduct,
instead of a monoadduct, was formed, allowing higher hydrogen storage
capacity of this system. The same catalyst also catalyzed the dehydrogenation
reaction. In the presence of 0.075 mol % Ru19, complete
dehydrogenation of a 1:1 adduct of HCOOH and DBU was observed at 100
°C in 70 min. Interestingly, the addition of toluene increased
the dehydrogenation rate. CO was not detected in the gas phase down
to 10 ppm. Under these catalytic conditions, up to five charge–discharge
cycles were performed in the same reactor, without deterioration of
the catalyst and high product yields (formate salts and H2 + CO2 mixture) were obtained at the end of each cycle.</p><p>Along the same direction, an LOHC based on HCOONa/CO2 in the basic medium was reported by Czaun, Prakash, Olah,
and co-workers.237 Both reactions, i.e.,
dehydrogenation of HCOONa
and hydrogenation of CO2 (in a basic medium), were catalyzed
by ruthenium pincer complexes under relatively mild conditions without
any need for external pH control. Dehydrogenation of formate salts
(20 mmol) was studied using pincer complexes Ru4, Ru1, or Ru1Me (20 mmol) in 1,4-dioxane/H2O solvent at 69–84 °C (Scheme 29). Both catalysts Ru4 and Ru1 exhibited similar catalytic activities (initial TOF =
286 h–1) for the dehydrogenation of HCOONa, and
an almost quantitative conversion of HCOONa was observed in ∼300
min at 70 °C producing approximately 490 mL of hydrogen gas.
Interestingly, complex Ru1Me was found to be more active
in this reaction and quantitative dehydrogenation of HCOONa was observed
in 4 h under the same conditions, giving an initial TOF of 430 h–1. This suggests that the N–H moiety does not
influence the rate of catalysis, in contrast to several other reports
using analogous systems where N–H plays an important role in
(de)hydrogenation reactions. The reverse reaction, i.e., hydrogenation
of CO2 to formate (H2/CO2 pressure
30/30 bar) in the presence of NaOH was also accomplished by using
complex Ru4 at 76 °C in THF/H2O solvent
resulting in the formation of sodium formate in 93% yield (TON = 1160).
Similar to dehydrogenation, complex Ru1Me was also active
for the hydrogenation reaction, suggesting that the N–H moiety
is not important in this transformation. Furthermore, hydrogenation
of bicarbonates and carbonates was also accomplished under similar
catalytic conditions.</p><p>Closed-loop cycles consisting of CO2 hydrogenation (charging)
and formate dehydrogenation (discharging) were also demonstrated by
using catalyst Ru4 at 78 °C up to six hydrogenation–dehydrogenation
cycles. Significantly, CO was not detected in the evolved gas mixture
in any of the catalytic cycles. Furthermore, the reversibility of
this hydrogen storage system was also demonstrated in the same pot.</p><p>On the basis of their mechanistic investigations, two possible
pathways for the catalytic cycle were proposed (Scheme 30). In both pathways, the first
step is the insertion of CO2 into the Ru–H bond
of complex Ru4a to form a ruthenium formate complex, Ru4b. In the first pathway (cycle I), the formate dissociates
to form the 16-electron complex Ru4c followed by its
reaction with H2 to produce the dihydrogen complex Ru 4d. The abstraction of a proton from complex Ru 4d results in the regeneration of the trans-dihydride
complex Ru4a and H2O. The second pathway (cycle
II) involves the formation of an amido complex, Ru4e,
by reaction of a base with Ru4b as proposed earlier for
the transfer hydrogenation of ketones using Ir-H3[(iPr2PCH2CH2)2NH].238 Addition of H2 to complex Ru4e regenerates the complex Ru4a. As cycle II involves deprotonation of the N–H
proton, catalysis performed by the complex Ru1Me occurs
most likely via cycle I (Scheme 30).</p><!><p>Reproduced with
permission
from ref (237). Copyright
2015 John Wiley and Sons.</p><!><p>Another challenge, in addition to storage, in front
of a viable hydrogen economy is that of sustainable production of
H2. As of 2018, ∼95% of hydrogen gas was produced
from fossil fuels, mainly via steam reforming of natural gas,239 resulting in dependence on the depleting fossil
fuels, and also producing large amounts of the greenhouse gas CO2. Relevant to this review are the opportunities that biomass
presents as a potential renewable feedstock for H2 production.240 Several catalysts, mostly heterogeneous, have
been reported for the production of hydrogen gas using gasification
or thermochemical conversion of biomass.240−242 Recently, a few transition metal complexes have been utilized for
the efficient production of H2 from biomass feedstock,
e.g., (bio)ethanol,243−245 glycerol,246−250 and polyols, e.g., sugar alcohols.251−254 Detailed review articles featuring dehydrogenation of alcohols catalyzed
by transition metal complexes have been reported recently.255−257</p><p>A very attractive strategy is water splitting, which offers
a sustainable way to produce clean H2 and O2, which is the reverse reaction occurring in a hydrogen fuel cell.258 Currently, water splitting is mainly performed
by electrolysis, but much research is directed at photoelectrochemical
water splitting and photocatalytic water splitting, as extensively
reviewed in recent years.259−265 An efficient and economic water-splitting technology can result
in a breakthrough needed to underpin the hydrogen economy. In a fundamentally
new approach, Milstein in 2009 reported a stoichiometric stepwise
water-splitting reaction where a ruthenium pincer complex splits H2O first thermally, producing H2 gas, and then photolytically
to produce O2 gas.266 The addition
of 1 equiv of H2O to the Ru-PNN dearomatized complex Ru13a resulted in the reversible O–H activation of
water via metal–ligand cooperation, forming the ruthenium hydride
hydroxo complex Ru20 (Scheme 31). Significantly, heating complex Ru20 in refluxing water for 3 days resulted in the formation
of the cis dihydroxo complex Ru20a with
concomitant evolution of hydrogen gas. Moreover, irradiation of complex Ru20a (in THF or water) with a 300 W halogen lamp filtered
through Perspex for 2 days resulted in the regeneration of the ruthenium
hydrido-hydroxo complex Ru20 with the concomitant formation
of oxygen gas. It was suggested that complex Ru20a eliminated
H2O2 via reductive elimination of two hydroxo
ligands followed by decomposition of H2O2 to
form H2O and O2. Labeling experiments were performed
to confirm that O–O bond formation in this system is intramolecular
and involves only a single metal center (Scheme 31).</p><p>Soon after this, successive DFT calculations were performed
by
Yoshizawa267 and Hall.268 Both calculations suggested that the heterolytic coupling
of ruthenium hydride with a proton from the side arm of the pincer
ligand in Ru20 forming the dearomatized sigma-complex Ru20c is the rate-limiting step. Further lower energy pathways
have been proposed independently by Suresh269 and Fabris,270 where solvolysis by H2O was found to play an important role. Suresh performed a
systematic study to explore how modification of Milstein's
RuPNN complex can affect the rate-determining H2 elimination
step.271 The study revealed that decreasing
the steric bulk of the phosphine moiety allows the exothermic association
of a water molecule with the ruthenium center and makes the intermediates
and transition states more stable. Thus, replacing the bulky tert-butyl group with a methyl or ethyl group at the phosphine
moiety of the pincer complex can reduce the ΔG⧧ by a significant amount, such as ∼10 kcal/mol.
As Ru–H facilitates H···H interaction with H2O for H2 elimination in the water-splitting reaction,269 Suresh studied the hydridic character of several
transition metal hydride complexes using the approach of the molecular
electrostatic potential (MESP).272 The
study suggested that a lower barrier for the activation of H2 can be obtained if a more electron-rich hydride ligand is used.
On the basis of DFT calculations, Fang273 proposed a new mechanism for the formation of triplet O2 that does not involve H2O2 intermediacy and
involves the formation of a dimer, Ru20S0 (Scheme 32). Photoexcitation
of this dimer results in a change of a spin state and the formation
of the Ru20T1 complex via the intersystem crossing (ISC)
of Ru20S1. This is followed by the first hydrogen atom
transfer (HAT) and the concerted dehydration steps to form Ru20d. Here the second HAT and the dehydration steps occur which are accompanied
by a concerted formation of the O–O bond in the complex Ru20e. The elimination of triplet O2 and H2O molecules from Ru20e results in the formation
of Ru20b. This mechanism agrees with the experimental
results from the Milstein group.</p><!><p>Our
current demand for transportation fuel is largely met by fossil fuels,
causing environmental and sustainability concerns. Biofuels which
can be produced from renewable biomass have been proposed as a sustainable
alternative to fossil fuels.274−278 One of the examples of biofuels is bioethanol that can be produced
via fermentation of biomass. However, direct use of ethanol as a fuel
has some drawbacks such as its low energy density, corrosive nature,
and formation of an azeotrope with water causing separation problems,
which limits its application in the transportation sector. On the
other hand, butanol isomers have a high energy density and noncorrosive
nature and are immiscible with water, more closely resembling gasoline.
Therefore, butanol isomers such as n-butanol are
termed "advanced biofuels".279−281 Thus, the synthesis of butanol from a bio-based feedstock such as
ethanol is an attractive topic. An important approach in this direction
is based on the "Guerbet reaction",282,283 discovered by Guerbet more than 100 years ago using simple sodium
alkoxides as catalysts at elevated (200 °C) temperature (Scheme 33A).284,285 Although simple and elegant, the process suffers from a selectivity
issue as a mixture of oligomers and polymers is formed under the reaction
conditions. As transition metal catalysts can allow the control of
kinetics and the oligomerization process, the "Guerbet reaction"
has been expanded to a "hydrogen borrowing reaction"
catalyzed by transition metal complexes. A proposed mechanism has
been outlined in Scheme 33B. The alcohol is first dehydrogenated to the corresponding
aldehyde, followed by self-aldol condensation of the formed aldehyde
to form an enal (crotonaldehyde) and subsequent hydrogenation of both
C=C and the aldehyde group of the enal (crotonaldehyde) to
form a higher-chain alcohol. Several heterogeneous286,287 and homogeneous catalysts have been reported in recent years for
catalytic upgradation of ethanol to butanol using a "hydrogen
borrowing" process. With relevance to the current report, we
discuss the details of homogeneous catalysts, in particular pincer
complexes, for ethanol to butanol transformation.</p><p>A homogeneous catalyst based
on iridium was reported by Ishii in
2009 for the upgradation of ethanol to butanol.288 Using EtONa in the presence of [Ir(cod)(acac)] (Ir8) (acac = acetylacetonate, cod = 1,5-cyclooctadiene), with 1,3-bis(diphenylphosphino)propane
(dppp) ligand and 1,7-octadiene as additive, n-butanol
was obtained with selectivity up to 61% but at a lower conversion
of 12%. Later, in 2013, the Wass group reported that using a ruthenium
complex, [trans-RuCl2(dppm)2] (Ru21) (dppm = 1,2-bis(diphenylphosphino)methane)
(0.01 mol %), and EtONa (5 mol %) at 150 °C and 20 h, 13% overall
conversion of ethanol was obtained with 90% selectivity to n-butanol (TON = 1330).289 The
same group also reported a more active system with in situ generated catalysts from a mixture of [RuCl2(η6-p-cymene)]2 and 2-(diphenylphosphino)ethylamine.
With the use of 0.1 mol % catalyst loading and EtONa base (5 mol %),
a good overall conversion of ethanol was obtained exhibiting high
selectivity (≥92%). The isolated active catalyst [RuCl(η6-p-cymene)(2-(diphenylphosphino)ethylamine)]Cl
(Ru22) showed almost the same reactivity.290</p><p>Around the same time, Mu reported that,
using [Ir(OAc)3] (Ir9) and bathophenanthroline
disulfonic acid disodium
salt as the ligand (L3), KOH, and NaOAc, resulted in
52% conversion of ethanol with 26% yield of n-butanol
at 150 °C for 16 h in air.291 Interestingly,
Jones demonstrated that, using highly basic, bulky complexes, the
nickel complex Ni1, or the copper complex Cu1 instead of bases such as NaOEt, excellent selectivity of butanol
can be obtained.292 The high yield/selectivity
was attributed to the sterically bulky nickel or copper hydroxide
complexes that catalyze the aldol reaction in almost quantitative
yield and selectivity to form the crotonaldehyde. The tandem catalytic
reaction was performed in a combination of the iridium complex (Ir10) used for the catalytic (de)hydrogenation steps and Ni1 or Cu1 complexes used for the catalytic aldol
reaction (Scheme 34).</p><!><p>TpMeMe = tris(3,5-dimethyl-pyrazolylborate).</p><!><p>Along this direction, the Szymczak group in 2016
demonstrated an
ethanol upgrade using an air-stable pincer-based complex Ru23 (Scheme 35).293 They reported that, by heating 17.1 mmol of
EtOH, 5 mol % NaOEt base, and 0.1 mol % Ru pincer catalyst (Ru23) at 150 °C for 2 h, 30% overall conversion of ethanol
to n-butanol (91% selectivity, with TON = 300) was
obtained. Significantly, performing the reaction with air-saturated
solvents and weighing all the reagents in the air did not affect the
reactivity of the system. The TON reached up to 1400 with 100% selectivity
at 1.4% conversion upon decreasing the catalyst loading to 0.001 mol
%. Detailed mechanistic studies suggested that displacement of PPh3 ligand and formation of mono- or dicarbonyl species slows
down the reaction rate which could be avoided by adding excess PPh3 ligand during catalysis.</p><p>A record TON for the ethanol-to-butanol transformation was reported
by Milstein in 2016 using an acridine-based ruthenium complex (Ru15, Scheme 36).294 A high yield and selectivity were
obtained at a temperature of 150 °C using EtONa as a base. Interestingly,
C6 and C8 alcohols, which have higher energy
capacities than butanol, were also formed as coproducts via cross-coupling
and homocoupling of 1-butanol. A maximum TON of 18 209 was
obtained after 7 days.</p><p>In addition to precious
metal complexes, a few base-metal catalysts
have also been recently reported for upgradation of ethanol to butanol.
The first example of a base-metal catalyst for this transformation
was reported by Liu in 2017 using a manganese pincer complex (Mn2) and NaOEt base (Scheme 37).295 Remarkably, a very
high catalytic activity was observed and by using 0.0001 mol % (8
ppm) Mn2, a record TON of 114 120 with 92% butanol
selectivity was obtained in 7 days (Scheme 37). Soon after, the Jones group also reported
the same transformation using the Mn2 complex (0.5 mol
%) and NaOEt base (25 mol %).296 A lower
amount of base loading (e.g., 50 equiv of EtONa relative to Mn2) was used in comparison to that reported by Liu (600 equiv
of EtONa relative to Mn2), which could be responsible
for a higher TON in Liu's case. Interestingly, both groups
reported a lower TON while using Mn2Me instead of Mn2, suggestive of the important role of N–H proton
in catalysis. A detailed theoretical investigation of this process
using a Mn(I)-PNP pincer complex has been reported by Pathak.297</p><p>Another attractive direction
for ethanol upgradation is the conversion
of ethanol to isobutanol, which has an improved octane number over
1-butanol. Ethanol can be coupled with methanol via a "hydrogen-borrowing"
pathway in the presence of transition metal catalysts to produce isobutanol
via a mechanism (Scheme 38) analogous to the ethanol-to-butanol transformation (Scheme 33B). As outlined
in Scheme 38, ethanol
and methanol first are dehydrogenated to form acetaldehyde and formaldehyde,
respectively, followed by their aldol coupling reaction to form acrylaldehyde
which subsequently is hydrogenated to form propanol. Propanol and
methanol undergo the same consecutive cycle to form isobutanol. A
few heterogeneous catalysts have been reported for this transformation,
but they suffer from limitations such as harsh reaction conditions
and low selectivity.298−301</p><p>A few homogeneous catalysts have also been investigated
for the
selective upgradation of ethanol to isobutanol. The first homogeneous
catalyst in this direction was reported by the Wass group in 2016
using [RuCl2(L)2]-type complexes in the presence
of a base such as NaOMe, where "L" is chelating diphosphine
or PN ligands.302 The best performance
was obtained by using the small-bite-angle dppm ligand based complex Ru21 which resulted in up to 75.2% conversion and almost quantitative
selectivity of isobutanol (Scheme 39). Later, the same group demonstrated that the Ru catalyst
(Ru21) remains active in the presence of water of a concentration
similar to that of fermentation broth.303 Interestingly, various commercial alcoholic beverages such as raki,
lager, gin, brandy, and wine were used as ethanol surrogates obtained
from fermentation broth instead of pure ethanol, exhibiting conversions
of up to 79%. Mansell, in 2017, also reported a ruthenium complex
bearing a chelating bis-phosphine ligand (Ru24) for the
coupling of methanol and ethanol to form isobutanol in 50% yield and
a selectivity of 96% (Scheme 39).304 Recently, in 2019, Xu, Yu,
and Liu reported the upgradation of ethanol to isobutanol using the
first base-metal catalyst based on manganese (Mn2, Scheme 39).305 Remarkably, a higher TON of up to 9233 (at
200 °C, 48 h) was obtained in comparison to those reported by
Wass (TON = 750 at 180 °C, 20 h) and Mansell (TON = 495 at 180
°C, 20 h).</p><!><p>Biodiesel is a type of diesel fuel compatible with
the current infrastructure of distributing transportation fuel. Typically,
biodiesel is a fatty acid methyl ester (FAME) that is usually produced
from the transesterification of triglycerides derived from animal
fat or plant oil with methanol.306−308 These FAMEs contain
unsaturated fatty acid derivatives such as linoleic acid and linolenic
acids. The direct use of such unsaturated FAMEs as biodiesels has
several drawbacks such as lower fuel stability and lubricity, increased
viscosity and gum formation, and slow ignition and high emission of
hydrocarbons. To avoid these drawbacks, hydrogenation of unsaturated
C=C has been performed using heterogeneous catalysts, under
harsh reaction conditions.309−312</p><p>Homogeneous catalysts have also been
employed lately to produce (saturated) biodiesel from vegetable oil.
The Williams group reported the transformation of a corn or soybean
oil to biodiesel and lactate in the presence of 0.3 mol % Ir11 and 5 equiv of NaOH in methanol or glycerol (Scheme 40).313 An advantage
of this methodology over the conventional biodiesel production technology
is that it produces lactate as a byproduct instead of glycerol, which
is discarded as a waste from the biodiesel industry. A satisfying
result was obtained using only 30 ppm Ir11 catalyst,
delivering over 230 000 turnovers in a single catalytic run.
Interestingly, by using a combination of Ir11 catalyst
and the iron pincer catalyst Fe3, a reduction of unsaturated
C=C was also obtained (product yield of up to 90%) without
loss of selectivity in conversion of the backbone to lactate (Scheme 40). Also, the system
exhibited good water tolerance, and the iridium catalyst was soluble
in water which allowed smooth reusability of the catalyst.</p><!><p>Lignin is
a biomass-derived complex polymer containing aryl alkyl ether and
alcoholic C–O linkages (Scheme 41). Lignin is usually considered a waste
product from the pulp and paper industry and is produced at the scale
of 70 million tons annually worldwide. Currently, lignin waste is
mostly used to generate energy by burning, leading to harmful impacts
on the environment.314 However, as lignin
is made of organic fragments, its depolymerization can provide a sustainable
route to access renewable and useful organic chemicals.</p><!><p>Important bonds for the purpose
of depolymerization are highlighted.</p><!><p>In recent
advances in this field, lignin has been depolymerized
into oligomers containing phenolic fragments, also called lignin oil,
which has applications in jet fuel and the polymer industry.315 However, the demonstration is only at the proof-of-concept
level and there are certain challenges associated with the depolymerization
of lignin on an industrial scale. These are due to (a) lack of a well-defined
polymeric structure, making it difficult to develop a general protocol,
and (b) difficulty in cleaving strong β-O-4 linkages (Scheme 41). Due to the need
for harsh reaction conditions such as a high temperature, a vast majority
of studies on the topic of lignin depolymerization involve heterogeneous
catalysts or acid/base catalysts.54−62 However, recently a few examples of well-defined transition metal
catalysts have been utilized for the reductive cleavage of the C–O
bond in model ether compounds of lignin (Scheme 42).</p><p>A breakthrough in this direction was reported by the Hartwig group
in 2011, who demonstrated selective cleavage of C–O bonds in
aryl ethers using soluble catalysts consisting of Ni(COD)2 (5–20 mol %), and N-heterocyclic carbene ligand (10–40
mol %), in the presence of NaOtBu (Scheme 42).316 Remarkably, only 1 atm H2 was sufficient
for the cleavage in the 80–120 °C temperature range. Soon
after, in 2012 and 2013, catalytic cleavage of cyclic alkyl ethers
was reported by Marks using Yb(OTf)3/Pd nanoparticles and
H2 through dehydroalkoxylation.317,318 Ellman and Bergman used the xantphos ligand with RuH2(CO)(PPh3)3 as a catalyst to cleave C–O
bonds in alkyl aryl ethers such as 2-phenoxy-1-phenethanol, where
the adjacent alcohol fragments transfer hydrogen to cleave the C–O
bond.319 In this direction, a ruthenium–xantphos
complex and a combination of [Ru(cod)(methylallyl)2] and
the triphos ligand was reported by James320 and Leitner321 respectively for the cleavage
of the same β-O-4 motif. All these catalysts used mild pressure
of H2 (1–40 bar) or an internal hydrogen donor moiety
at temperatures in the range 120–185 °C and produced moderate
to excellent yields of the products.</p><p>A new atom-economical approach
without using any reductant was
reported by Goldman. This was based on dehydroaryloxylation of alkyl
aryl ethers using an Ir-based pincer catalyst.322 Various aryl ethers were successfully transformed to phenol
and an olefin in the presence of 2 mol % (iPrPCOP)Ir (Ir12) at 150–200 °C for 16 h.
Based on stoichiometric control experiments, a mechanism for the dehydroaryloxylation
reaction was proposed as outlined in Scheme 43. The first step is the C–H activation
of the alkyl aryl ether to form a four-membered cyclometalated complex
(Ir12a). This is followed by the migration of the aryl
oxide moiety forming an olefin-bound complex Ir12b that
liberates olefin to form Ir12c followed by the reductive
elimination of phenol, regenerating Ir12. Direct depolymerization
of lignin has also been attempted using a homogeneous ruthenium/xantphos
catalyst; however, the catalytic activity was poor.323</p><!><p>Methanol is a highly important feedstock for several
high-value
chemicals and is globally produced at the scale of more than 75 million
metric tons.324 Additionally, methanol
has significant applications in the energy sector in internal combustion
engines (ICE) and in direct methanol fuel cells (DMFC). The application
of methanol as a hydrogen carrier for the hydrogen economy and homogeneous
catalysts developed for the aqueous methanol reforming reaction has
been discussed in section 2.1. Olah and Prakash have proposed the vision of a "methanol
economy",26,64,325,326 where methanol can be essentially
produced in a renewable, sustainable, and carbon-neutral cycle by
the carbon capture and recycling (CCR) process. In this section, we
discuss homogeneous catalytic processes that have been developed for
the sustainable production of methanol.</p><!><p>Currently, methanol is industrially
produced from syngas327 that is produced
from fossil fuels, such as
the gasification of coal.328 However, recently
it has been possible to produce methanol from biomass or by the direct
hydrogenation of CO2.329 Industrial
production of 100% renewable methanol is carried out by Carbon Recycling
International at the scale of 50 000–100 000
tons/year, where CO2 is captured from industrial emissions
and then hydrogenated to methanol by using H2 produced
from the electrolysis of water using renewable electricity.330 Most of the reports on direct hydrogenation
of CO2 to methanol (CO2 + 3H2 →
CH3OH + H2O) involve heterogeneous catalysts.331−333 Electrochemical reduction of CO2 to methanol using molecular
complexes has also received significant attention.334−337</p><p>Hydrogenation of CO2 to methanol is exothermic
in nature (ΔHr,298 K = −50
kJ/mol); however, high temperature is required to overcome the chemical
inertness of CO2 that poses a thermodynamic hindrance as
the reaction becomes thermodynamically less favorable at higher temperature.
To overcome the thermodynamic challenge, an indirect approach for
the conversion of CO2 to methanol has been developed using
homogeneous catalysts under relatively mild conditions.338−343 In this approach, CO2 is first chemically captured by
its reaction with a trapping reagent such as amines, alcohols, silanes,
or boranes, followed by subsequent hydrogenation or hydrolysis to
form methanol and regenerate the trapping reagent (Scheme 44). Section 4.1.1 describes various transition metal catalysts
used for the indirect conversion of CO2 to methanol.</p><!><p>Conventionally amines are used to capture
CO2 as a carbamate salt from which CO2 can be
liberated at a high temperature with concomitant regeneration of amines.
However, a significant amount of amine is decomposed in the regeneration
process, leading to the low efficiency of this process. As amine-assisted
CO2 capture technology has already been developed, transformation
of the trapped intermediate to methanol can be an attractive indirect
approach for CO2 reduction to methanol. Amine-assisted
CO2-to-methanol transformation was first reported by Sanford
using NHMe2 and a pincer catalyst, Ru4 (Scheme 45).344 CO2 is captured using NHMe2 to form dimethylammonium dimethylcarbamate (DMC) that can be subsequently
hydrogenated by Ru4 to form a mixture of methanol and
dimethylformamide (DMF). With the use of a combination of Ru catalyst
(Ru4) and K3PO4, CO2 was hydrogenated to a mixture of DMF/DMFA (dimethylammonium formate)
and CH3OH. Overall, 96% conversion of CO2 and
27% yield of methanol were obtained.</p><p>Around the same time,
the Milstein group reported hydrogenation
of low-pressure CO2 (1 atm).345 CO2 was captured using valinol to form an oxazolidinone-type
intermediate followed by subsequent hydrogenation to methanol. CO2 capture was catalyzed by Cs2CO3,346 whereas the hydrogenation reaction was catalyzed
by a ruthenium pincer complex Ru6 in the presence of
KOtBu (Scheme 46). Furthermore, a tandem reaction was also
demonstrated without the need for the isolation of the oxazolidinone
derivative. Only 1 atm pressure of CO2 was sufficient for
the formation of oxazolidinone derivative which was hydrogenated to
form methanol at 60 bar and 135 °C, producing MeOH in 53% yield
(Scheme 46).</p><p>In a similar direction, Ding
reported a sequential CO2 reduction to CH3OH
using a Ru-MACHO catalyst, Ru1, in the presence of morpholine.347 Under
the reaction condition (70 atm of 1:1 CO2:H2) CO2 was captured as N-formylmorpholine,
which was subsequently hydrogenated to form methanol in 36% yield
(Scheme 47).</p><p>Prakash has recently reported a series of papers on CO2-to-methanol transformation using various additives. For example,
in an interesting discovery, CO2 was hydrogenated to CH3OH in the presence of polyamines such as pentaethylenehexamine
(PEHA; Scheme 48).348 Polyamines have higher basicities, allowing
them to capture more CO2 than an amine molecule. Moreover,
they exhibit higher thermal stabilities and lower volatilities, making
them better candidates for CO2 capture. The ruthenium-MACHO-BH
catalyst Ru4 was utilized for the hydrogenation of CO2 at the pressure of 75 bar (CO2/H2 1:3)
and in the temperature range 95–155 °C in THF solvent.
The addition of a catalytic amount of base (K3PO4) was found to increase the yield of methanol from 7.6 to 9 mmol
presumably by favoring the N–H-assisted metal–ligand
cooperation pathways. Recycling of catalyst was also demonstrated
for five cycles exhibiting a total TON of 1850. A proposed mechanism
that is similar to those using amine for CO2 capture is
outlined in Scheme 48. The reaction starts with the trapping of CO2 by basic
polyamines resulting in the formation of a carbamate salt that is
hydrogenated to form a formate salt. The formate salt is dehydrated
under the thermal condition to form a formamide that subsequently
is hydrogenated in the presence of Ru4 to form methanol
and regenerate the polyamine.</p><p>Later, Prakash reported a more
practical approach of amine-assisted
CO2-to-methanol transformation using a biphasic solvent
system consisting of 2-MeTHF and H2O, considering that
catalysts are mostly soluble in organic solvents whereas the capturing
amines are soluble in water.349 This approach
was helpful in the recycling of amines and catalysts after the hydrogenation
step. Mechanistic studies were later reported by Prakash.350 Interesting structure–activity studies
were conducted by the variation of phosphine substituents of the Ru-PNP
complex, revealing that methanol yield followed the order Ph > iPr > Cy > tBu
(phosphine substituents, PR2). Based on further experiments,
a mechanism consisting of two cycles as outlined in Scheme 49 was proposed. Cycle A describes
the formation of formamide from CO2, amine, and H2, whereas cycle B describes the hydrogenation of formamide to form
methanol. The catalysis starts with cycle A where the dihydrogen sigma-complex Ru1a, formed from the reaction of precatalyst Ru1 with base and H2, is deprotonated by an amine, forming
a ruthenium trans-dihydride complex Ru1b. Insertion of CO2 to Ru1b results in the
formation of the formyl complex Ru1c. Attack of an amine
on the formyl moiety of Ru1c forms an alkylammonium formate
salt and the cationic complex Ru1d that coordinates with
H2 to regenerate the sigma-complex Ru1a. The
formed alkylammonium formate salt is dehydrated under thermal conditions
to form formamide that feeds to cycle B, where it is hydrogenated
by Ru1b to form a hemiaminal intermediate and the complex Ru1e. Ru1e reacts with H2 to regenerate
the trans-dihydride complex Ru1b. The
hemiaminal intermediate decomposes to form formaldehyde that is hydrogenated
by the same cycle to form methanol.</p><!><p>P = PPh2.</p><!><p>Furthermore, Prakash reported a modified version
of this process
by immobilizing the amine onto a solid support.351 The advantage of solid-supported amines (SSAs) is that
no solvent is needed, making the overall process less energy intensive.
A nonpincer ruthenium complex Ru25 was utilized by Wass
for the hydrogenation of CO2 to methanol in the presence
of an amine (Scheme 50).352 Similar to the mechanism described
earlier, CO2 is trapped by an amine and H2 to
generate a formamide that is hydrogenated by Ru25 to
form methanol, regenerating the amine. It was observed that the choice
of amine plays a crucial role in the yield of methanol. Bulkier amines
such as pyrrolidine, iPr2NH,
and Et2NH afforded lower yields compared to that of Me2NH. No methanol was produced by using the tertiary amine NEt3 due to the lack of a N–H proton required for the generation
of a formamide. Catalysis was performed at 180 °C, under 10 bar
CO2 and 30 bar H2 pressure, exhibiting a TON
up to 8900 and a TOF up to 4500. Interestingly, no methanol was formed
when the catalysis was performed using an analogous ruthenium complex
containing a tertiary amine moiety, suggesting the crucial role of
the N–H proton of the ligand in catalysis. It was suggested
that the catalytic hydrogenation occurs via an "outer-sphere
mechanism" facilitated by metal–ligand cooperativity.</p><p>In addition to precious metal
complexes, base-metal complexes were
also employed for the integrated conversion of CO2 to CH3OH (Scheme 51). The Prakash group employed a manganese pincer catalyst, Mn2, for the conversion of CO2 to CH3OH using amines such as morpholine or benzylamine (Scheme 51).353 The CO2 is trapped as an N-formylated product that is
hydrogenated to produce methanol. Compared to the ruthenium catalysts
described above, the TON was found to be significantly lower (up to
36). Using a similar approach, the Bernskoetter group recently reported
the hydrogenation of CO2 to methanol catalyzed by an iron
pincer catalyst, Fe2, using morpholine as an amine source
(Scheme 51).354 The initial step involves N-formylation of
morpholine using CO2/H2 to produce formylated
morpholine and is followed by deaminative hydrogenation in the presence
of an iron pincer catalyst to form methanol. Overall, a high TON of
1160 was obtained for the reaction of CO2 and H2 to form formylmorpholine, and a TON of 590 was achieved for the
hydrogenation of isolated formylmorpholine to methanol. Mechanistic
studies suggest that the presence of CO2 inhibits the hydrogenation
of formamide by forming a stable iron(II) formate complex, thus hampering
the single batch catalysis. Along this direction, Martins and Pombeiro
have reported an iron catalyst exhibiting a TON up to 2300 for the
hydrogenation of CO2 to methanol using pentaethylenehexamine.355 Drawing inspiration from the metal–ligand
cooperation exhibited by [FeFe] hydrogenase enzymes, the authors used
the tripodal C-scorpionate iron(II) complex (Fe4) as
a catalyst (Scheme 51). It is expected that the N-atoms present in the pyrazolyl rings
assist in proton transfer needed for the H–H cleavage step.</p><!><p>In addition to amines,
alcohols have also been employed for the
reduction of CO2 to CH3OH (Scheme 52). The first example in this
direction was reported in 2011 by the Sanford group using a cascade
catalytic process.356 The cascade catalysis
consisted of (a) hydrogenation of CO2 to HCOOH catalyzed
by Ru26, (b) Sc(OTf)3 catalyzed reaction of
HCOOH with methanol to form methyl formate (HCOOMe), and (c) hydrogenation
of HCOOMe catalyzed by Milstein's catalyst (PNN)Ru-(CO)(H)
(Ru6a) to form methanol. To demonstrate the proof of
concept, a reaction of 30 bar CO2 and 10 bar H2 was performed using 0.0126 mmol each of the catalysts Ru26, Sc(OTf)3 (Sc1), and Ru6a,
which resulted in the formation of a mixture of CH3OH (2.5
turnovers) and HCO2CH3 (34 turnovers). The low
yield of CH3OH was attributed to the deactivation of Milstein's
catalyst (Ru6a) by Sc(OTf)3. This issue was
solved by physically separating the Ru6a catalyst from
the remaining two catalysts, which were stored in a vial in the central
area of the reactor, while the PNN complex Ru6a was stored
in the outer wall of the autoclave. The in situ generated
methyl formate that formed in the central vial was transferred to
the outer vessel under the reaction condition and was hydrogenated
in the presence of Ru6a. This approach was indeed successful,
and a higher yield of methanol (21 turnovers) was obtained.</p><p>Goldberg has recently reported
a more efficient catalytic system
using a similar approach.357 Screening
of several catalysts revealed that the combination of Ru27/Sc1/Ir13 was the most active for the hydrogenation
of CO2 to methanol in the presence of ethanol. A TON of
428 (40 h, 155 °C) was obtained under the catalytic condition
producing a high yield of methanol (1.07 M).</p><p>Byers and Tsung
have recently reported a bio-inspired multicomponent
catalytic approach for the hydrogenation of CO2 to methanol
in the presence of alcohol.358 The cascade
catalytic process involves two catalysts: (a) a RuPNP pincer catalyst
encapsulated in an MOF (UiO-66) that converts CO2 to a
formate ester via HCOOH in the presence of alcohol and (b) Milstein's
catalyst Ru6 (Table 2) that hydrogenates the formate ester to methanol and
regenerates the alcohol. The heterogeneous catalyst system was successfully
recycled, demonstrating an excellent TON of up to 21 000 at
the end of the five cycles. The same groups have recently studied
the effect of the secondary sphere interaction between the encapsulated
pincer catalyst and the MOF host on the catalytic hydrogenation reaction.359 A variety of functionalized MOF (UiO-66-X)
hosts was used to study the structure–activity relationship.
The best results were obtained using the ammonium functional group
(UiO-66-NH3+) that resulted in a significantly
higher TON of up to 19 000. The catalyst recyclability was
also demonstrated with a TON of up to 100 000 at the end of
10 cycles. The high activity was attributed to the enhanced rate of
the hydrogenation of CO2 to HCOOH which was accelerated
due to the presence of ammonium moiety.</p><p>In addition to pincer
catalysts, systems based on triphos ligands
have also been studied for the conversion of CO2 to CH3OH in the presence of alcohol additives. A key feature of
such systems is that they operate under acidic conditions, making
them distinct from most of the pincer catalysts that operate under
basic or neutral conditions. Klankermayer, Leitner, and co-workers,
in 2012, reported the ruthenium-triphos catalyst Ru28 in combination with the acid cocatalyst bis(trifluoromethane)sulfonimide
(HNTf2) for the efficient hydrogenation of CO2 to CH3OH (Scheme 53).360 Ethanol was used as an additive to trap the formic acid as ethyl
formate that could be easily hydrogenated. Mechanistic studies reported
later by the same group revealed that HNTf2 reacts with Ru28 to form a cationic species Ru28a that was
proposed to be an active species in the catalytic process.361</p><p>In addition to direct hydrogenation, transfer hydrogenation has
also been reported recently by Klankermayer using a combination of
[Ru(triphosXyl)(tmm)] (Ru29, Scheme 53) and the Lewis acid Zn(NTf2)2.362 Linear alcohols such as ethanol
were used as the hydrogen source as well as to stabilize formic acid
by forming ethyl formate and a TON of up to 121 was achieved under
relatively mild conditions. Very recently the same group also demonstrated
a highly active cyclohexyltriphosphine ligand based ruthenium catalyst
(Ru30, Scheme 53) system for the hydrogenation of CO2 to methanol
using aluminum tristriflate (Al(OTf)3) as a Lewis acid
additive.363 Under the optimized conditions
of H2/CO2 = 90/30 bar pressure, 120 °C,
and 20 h reaction time, a TON of up to 2100 was obtained in ethanol
solvent. Moreover, the reduction of CO2 to methanol was
also performed in a biphasic mixture consisting of n-decanol and water with the same system under the optimized reaction
condition with maximum TONs up to 1087.</p><p>In addition to ruthenium-triphos
systems, base-metal triphos catalysts
were also utilized for the conversion of CO2 to methanol.
Beller's group utilized the Co(acac)3/triphos-based
system (Co1, Scheme 53) in the presence of HNTf2 as an additive
for the hydrogenation of CO2 to methanol.364 A TON of up to 50 was obtained under the condition of 70
bar H2 and 20 bar CO2 at 100 °C. Later
the same group also reported a more effective cobalt catalyst for
the hydrogenation of CO2 to CH3OH using a cobalt
complex based on a modified triphos ligand (Co2, Scheme 53).365 The p-toluene substituted
triphos ligand system with Co(acac)2 and HNTf2 as an additive led to TONs up to 125. Remarkably, the system also
worked efficiently under additive-free conditions by replacing the
metal precursor with Co(NTf2)2.</p><p>In addition
to methanol and ethanol, alkali metal hydroxide solutions
have also been utilized for the integrated capture and transformation
of CO2 to methanol. Alkali metal hydroxide solutions have
several advantages such as higher abundance, higher stability, and
higher efficiency for CO2 capture from the air in comparison
to amines.366</p><p>Recently, Prakash reported
the use of a solution of an alkali metal
hydroxide such as KOH for CO2 capture, and ethylene glycol
as formate ester stabilizer, to convert CO2 to CH3OH catalyzed by a ruthenium pincer catalyst, Ru4.367 CO2 was captured as carbonate and
alkyl carbonate salt by bubbling air through the KOH solution. This
was treated with 0.5 mol % (Ru4) catalyst and 70 bar
H2 at 140 °C, which resulted in a methanol yield of
25% within 20 h and a quantitative yield after 72 h. A mechanistic
pathway was speculated based on catalytic experiments (Scheme 54). The metal alkyl carbonate
salt produced from the reaction of CO2, ethylene glycol,
and KOH reacts with H2 to form ethylene glycol and potassium
formate salt. HCOOK reacts with the generated ethylene glycol to form
2-hydroxyethyl formate that is subsequently hydrogenated to methanol,
regenerating ethylene glycol.</p><p>Recently, Prakash has expanded this concept and reported
a tertiary
amine–ethylene glycol based integrated CO2 capture
and hydrogenation process.368 The combination
of a tertiary amine and ethylene glycol was found to efficiently capture
CO2 to form an ammonium alkyl carbonate intermediate that
subsequently is hydrogenated in the presence of Ru4 catalyst
to produce methanol. Of several screened tertiary amines in combination
with ethylene glycol, tetramethylethylenediamine and tetramethylbutanediamine
were found to afford the best yields of methanol (up to 94% using
0.5 mol % Ru4) upon CO2 capture/hydrogenation
from a gas mixture containing 10% CO2, similar to that
found in the flue gas. Although the proof of concept was successfully
demonstrated, the TON was similar to that of the earlier case where
ethylene glycol in combination with KOH was used and CO2 could be directly captured from air (Scheme 54).</p><!><p>In addition to
molecular hydrogen, reducing agents such as silanes
and boranes have also been employed for the transformation of CO2 to CH3OH. The silylated or boranated products,
formed from the reaction of CO2 with silanes or boranes,
respectively, can be easily hydrolyzed to form methanol. Their advantages
compared to hydrogenation reactions are that (a) they are easy to
handle, (b) reactions can be performed under mild conditions, and
(c) the reaction is thermodynamically more favorable because of the
formation of strong Si–O or B–O bonds. Catalysts based
on main group369−371 and transition metals372−377 have been studied in the past for the conversion of CO2 to methoxysilyl derivatives. Regarding transition metal pincer complexes,
the Chirik group reported a cobalt pincer catalyst, Co3, for the hydrosilylation of CO2 using PhSiH3 that produced a mixture of oligomers (Scheme 55).378 Along this
line, the Huang group reported the reduction of CO2 (1
atm) in DMF using Ph2SiH2 as a reductant and
the dearomatized PN3P*-Ni hydride complex Ni2 as a catalyst
(Scheme 55).379 Methanol was obtained in excellent yield (e.g.,
91%) from the hydrolysis of hydrosilylation product.</p><p>Similarly, Abu-Omar reported a rhenium pincer catalyst Re1 for the reduction of CO2 (100 psi) using Me2PhSiH (Scheme 56).380 The reaction initially produced
silyl formate
HC(O)(OSiMe2Ph), which was further reduced by the addition
of a primary silane, PhSiH3. Methoxysilane was obtained
in 53% yield from silyl formate, via the silyl formal intermediate,
by treatment of excess PhSiH3 for 24 h using catalyst Re1 (Scheme 56).</p><p>Along the direction of base-metal
catalysis, Kirchner and Gonsalvi
reported manganese catalyzed hydrosilylation of CO2 to
methoxysilane under mild conditions (80 °C, 1 bar CO2).381 Catalysis was performed with [Mn(iPrPN3P)(CO)2H] (Mn6) at room temperature, resulting in the rapid formation
of silylformates, which convert to methoxysilyl products over time
(93% in 46 h), whereas performing the reaction at a higher temperature,
e.g., 80 °C, shifts the selectivity to methoxysilyl products
(99% in 6 h). A mechanism of this reaction as elucidated experimentally
and by DFT calculations is outlined in Scheme 57. First, Mn–H attacks CO2 to form a metal-coordinated formate intermediate, Mn6a, which undergoes a reaction with silane to form the silyl formate
Mn hydride complex Mn6b. The formation of Mn6b from Mn6a has the highest barrier in the entire cycle.
Attack of the hydride ligand to the C=O bond of the attached
formate moiety in complex Mn6b forms a coordinated silyl
hemiacetal, Mn6c. This is followed by the attack of silicon
from a second PhSiH3 molecule to the OSiH2Ph
moiety of complex Mn6c resulting in the release of (PhSiH2)2O, forming the methoxy complex Mn6d. Finally, attack of the third silane molecule to the coordinated
methoxy ligand results in the formation of the final product, methoxyphenylsilane
((CH3O)SiH2Ph), with the concomitant regeneration
of the manganese hydride complex Mn6 (Scheme 57).</p><p>In addition to silanes, boranes were also used for CO2 capture via catalytic hydroboration followed by hydrolysis to methanol.
Guan used Pd382 and Ni383−385 complexes of the POCOP pincer ligand, which catalyzed hydroboration
of CO2 under atmospheric pressure to form a hydroborated
product which was subsequently hydrolyzed to form methanol (Scheme 58). The Pd-based
catalyst Pd1 was found to be air-stable and hydroborated
CO2 with catecholborane (HBcat) to form CH3OBcat
and catBOBcat at room temperature with TOFs up to 1780 h–1.382 The nickel-based pincer complex Ni3 catalyzed the same transformation at room temperature
and an atmospheric pressure of CO2, although with a lower
TOF of 495 h–1.383 The
catalytic activity of analogous pincer complexes containing different
PR2 (R = cyclopentyl, iPr,
and tBu) groups at the Ni3 complex were also studied, revealing that a more bulky substituent
at the phosphine moiety favors the catalytic hydroboration reaction.384 Moreover, the same group also reported a nickel
borohydride complex, Ni4, that catalyzed the hydroboration
of CO2 at 1 atm pressure and 60 °C using 9-BBN (Scheme 58).385</p><!><p>Although
synthesis of CH3OH directly from CO2 is more
sustainable and attractive, several studies have also been carried
out on the hydrogenation of CO to CH3OH. Hydrogenation
of CO to CH3OH is more exothermic (ΔH298 K = −90.7 kJ/mol) than that of CO2 to CH3OH (ΔH298 K = −49.5 kJ/mol), which could allow the process to occur under
relatively mild conditions. Seminal studies on the direct hydrogenation
of CO to CH3OH using homogeneous catalysts were reported
at harsh conditions.386,387 Recently, homogeneous catalysts
have been employed for the indirect hydrogenation of CO to CH3OH in the presence of an additive such as alcohol or amine.
This allows the process to occur under relatively mild conditions.
For example, Mahajan388 and Jens389,390 independently demonstrated hydrogenation of CO to CH3OH via methyl formate (at 90–140 °C by Mahajan and at
60–120 °C by Jens). The reaction proceeds by insertion
of CO into CH3OH to form methyl formate followed by hydrogenation
to form CH3OH (Scheme 59).</p><p>Along this line, amines were also used to capture CO to facilitate
the hydrogenation process. Prakash recently reported the hydrogenation
of CO to CH3OH by a two-step process: (1) trapping of CO
with an amine such as piperidine to form formamide using K3PO4 catalyst and (2) hydrogenation of the formamide intermediate
to CH3OH and piperidine catalyzed by the ruthenium pincer
complex Ru4 (Scheme 60).391 The yield of methanol
increased when a polyamine diethylenetriamine was used instead of
piperidine. Furthermore, direct hydrogenation of CO to CH3OH using diethylenetriamine and a catalytic combination of K3PO4 and the ruthenium catalyst Ru4 was also demonstrated, exhibiting a TON up to 570 (Scheme 60). This process has several
advantages, such as relatively low reaction temperature, absence of
corrosive alkali metal alkoxide bases, inexpensive K3PO4, and inexpensive high-boiling polyamines.</p><p>An analogous transformation
was also reported by the Beller group using a base-metal catalyst, Mn2, in combination with K3PO4 and an
amine promoter (Scheme 61a).392 The type of amine was found
to affect the yield of methanol. For example, TON > 500 was used
in
the case of indole, whereas scatole and pyrrole resulted in relatively
high TONs of >3000 and 2500, respectively. Recently the Leitner
group
reported the alcohol assisted CO reduction to methanol using the Mn
pincer complex (Mn2) resulting in a TON of 4023 and a
TOF of 857 h–1 in EtOH/toluene as solvent under p(CO/H2) = 5/50 bar at 150 °C (Scheme 61b).393</p><!><p>The
disproportionation of formic acid (FA) to MeOH and CO2 is
also a promising method for the indirect conversion of CO2 to MeOH, especially as the area of hydrogenation of CO2 to formic acid/formate has been significantly developed in
the past two decades (Scheme 62). The disproportionation reaction was first reported by Goldberg
and co-workers in 2013 using the iridium catalyst [Cp*Ir(bpy)(H2O)](OTf)2 (Ir13).394 A TON up to 156 was observed; however, a low methanol selectivity
(12%) was obtained. This was followed by a report from the Cantat
group that utilized the ruthenium-based catalyst Ru31 to produce methanol in 50.2% yield from the disproportionation reaction
of FA.395 Laurenczy and Himeda demonstrated
the same transformation with yields of up to 75% in D2O
catalyzed by an iridium complex (Ir14), resulting in
a TON of 240 (TON up to 1260 at the end of five cycles).396 Moreover, the same groups also reported a selective
disproportionation of formic acid to methanol with a selectivity of
96% using the same catalyst (Ir14) under isochoric and
acidic conditions.397 In 2016, the Kubiak
group demonstrated the electronic effects on the catalytic disproportionation
of HCOOH to CH3OH using cationic iridium bipyridine complexes
(Ir15, Ir16) and noticed that the unsubstituted
bipyridine complex and the 4-tBu substituted
complex exhibited the highest selectivity toward methanol.398 Soon after, Laurenczy and Himeda reported their
study on various iridium catalysts with substituted 2,2′-bipyridine
derivatives for disproportionation of formic acid. Their report revealed
that the iridium catalyst bearing 5,5′-dimethyl-2,2′-
bipyridine (Ir14a) showed high TON and selectivity toward
methanol.399 Detailed DFT studies were
also performed by Yang on this system.400 Furthermore, Himeda and Inagaki reported the catalytic disproportionation
of HCOOH catalyzed by heterogeneous catalyst Ir17 and
homogeneous Ir-bipyridine catalysts. The analogous heterogeneous catalyst Ir17a was found to be more selective (up to 8.3%) for methanol
production in comparison to the homogeneous Ir-bipyridine catalysts
(1.4–4.3%).401 Neary and Parkin
reported that the molybdenum complex CpRMo(PMe3)3–x(CO)xH afforded 21% selectivity of MeOH and methyl formate formation
in benzene at 100 °C.402 In 2016,
the Cantat group also reported a metal-free system for the formic
acid disproportionation reaction using stoichiometric quantities of
dialkylborane reagents (B01), where boroformate and borohydride
intermediates were formed via the decarboxylation of formate and followed
by undergoing disproportionation of formates to formaldehyde and methanol.403</p><!><p>PMO = periodic mesoporous
organosilica. Sol = solvent, e.g., acetonitrile.</p><!><p>Methane,
which constitutes 70–90% of natural gas,26 is a greenhouse gas having a 25 times larger
impact than CO2 on global warming.404 As long-distance transport of natural gas (in gaseous form)
is challenging, methane is flared to CO2 in oil fields,
converting a potential resource to waste. Thus, sustainable methods
for the conversion of CH4 to liquid fuels can play an important
role in meeting the energy demands of the future. Current approaches
for the transformation of CH4 to liquid fuels involve the
transformation of CH4 to gasoline or syngas using the steam
reforming process.405 However, both processes
require harsh reaction conditions. Another approach that has been
of high interest is partial oxidation of methane to methanol (CH4 + 1/2O2 → CH3OH), benefiting
the methanol economy. Although the reaction is exothermic in nature
(ΔH°298 K = −126.2
kJ/mol), the process is highly challenging due to the associated kinetic
barriers: (a) activation of the C–H bond in CH4 requires
harsh reaction conditions due to the high dissociation energy of the
C–H bond (440 kJ/mol) and (b) methanol C–H bond dissociation
is 47 kJ/mol lower than that of CH4, making methanol susceptible
for further oxidation under the reaction conditions. Thus, a suitable
catalyst is required that can activate the methane C–H bond
toward oxidation to methanol without overoxidation. Several heterogeneous
catalysts have been reported for the partial oxidation of CH4 to CH3OH, although most of them work under high temperatures.406 A few homogeneous catalysts have also been
reported for the oxidation of CH4 to CH3OH.
A common challenge in the direct oxidation of CH4 to CH3OH is low yield and selectivity due to the kinetic limitations
of this reaction. To address this issue, homogeneous catalysts have
been employed to convert methane to methyl esters that can be hydrolyzed
to produce methanol. A breakthrough in this direction was reported
by Periana, who reported the conversion of methane to methanol via
methyl bisulfate using a mercuric bisulfate catalyst.407 The catalysis starts with the C–H bond
activation of methane by mercuric bisulfate (Hg(OSO3H)2) catalyst to produce an observable species, CH3HgOSO3H, eliminating H2SO4. This
is followed by the decomposition of CH3HgOSO3H to CH3OSO3H and Hg22+. Hg22+ subsequently is oxidized by sulfuric
acid to regenerate the active species Hg(OSO3H)2 and produce byproducts SO2 and H2O (Scheme 63). The same group
later reported a catalyst involving a platinum bipyrimidine complex
in H2SO4 that afforded a higher yield of methyl
bisulfate (72%) compared to mercuric bisulfate catalyst that produced
methyl bisulfate in 45% yield. The mechanism is similar to that of
mercuric bisulfate catalyst as outlined in Scheme 63. After these seminal discoveries, several
other groups have reported the oxidation of CH4 to methanol
derivatives using homogeneous catalysts as reviewed recently.408−410 However, despite a good effort in the development of both heterogeneous
and homogeneous catalysts for the direct oxidation of methane to methanol,
there has been no commercial process in this direction. This is mainly
because of the higher reactivity of methanol than methane and the
thermodynamic formation of CO2 that leads to low yield
(concentration) and poor selectivity of methanol, creating a need
for the use of a stoichiometric additive.</p><!><p>Lower
alkanes can be derived from several sources, e.g., from biomass411 or by the reduction of CO2.412 Thus, alkane upgradation presents new opportunities
to produce conventional petroleum-based fuel such as diesel and jet
fuel (C9–C16) from renewable feedstocks (lower alkanes). However,
the inertness of C–H bonds in unactivated alkanes makes this
transformation challenging. Recent discoveries in catalysis have led
to two new approaches for alkane upgradation: alkane metathesis and
alkane–alkene coupling, as discussed below.</p><!><p>Due to the requirement
of high temperature to activate the inert C–H bonds of alkanes,
alkane metathesis was first reported using heterogeneous catalysis.
Burnett and Hughes in 1973 reported the disproportionation of alkanes
over a catalytic combination of platinum on alumina mixed with tungsten
oxide on silica.413 This was followed by
a series of papers from the Basset group describing the metathesis
reaction of linear or branched alkanes producing the next higher and
lower alkanes catalyzed by the silica-supported transition metal hydrides
(based on tantalum, chromium, or tungsten) at moderate temperatures
of 25–200 °C. An informative review of the earlier work
on alkane metathesis was reported by Basset.414</p><p>The first homogeneous catalytic system for alkane metathesis
was discovered by Goldman, Brookhart, and co-workers in 2006.415 Using a tandem combination of two independent
catalysts, an iridium pincer catalyst for alkane (de)hydrogenation
and Schrock's Mo1 catalyst for olefin metathesis,
efficient and selective metathesis of linear alkanes was achieved
at a moderate temperature. The alkane metathesis proceeds by first
alkane dehydrogenation to olefins followed by olefin metathesis to
form one higher-chain and another lower-chain olefin molecules. The
generated olefins are then hydrogenated by H2 formed in
the dehydrogenation step (Scheme 64). For example, heating an n-hexane
solution containing Ir18 and Mo1 at 125
°C in a sealed system under argon resulted in the formation of
a range of C2–C15 n-alkanes in 24 h. Interestingly,
no branched or cyclic alkanes were detected, unlike that reported
by Basset, Copéret, and co-workers.414 A combination of the iridium pincer catalyst (MeO-iPrPCP)IrH4 with the metathesis catalyst Re2O7/Al2O3 showed a significant improvement
in turnover numbers.</p><!><p>Reproduced from ref (416). Copyright 2017 American
Chemical Society.</p><!><p>Guan and Huang utilized
the concept of tandem, catalytic cross
alkane metathesis (CAM) for the degradation of polyethylenes (PEs)
into liquid fuels and waxes.417 A mixture
of polyethylene and a light alkane was heated under argon in the presence
of a combination of the iridium pincer complex Ir15 used
for alkane dehydrogenation and Re2O7/Al2O3 catalyst for the purpose of olefin metathesis.
As depicted in Scheme 65, the overall process occurs in three steps. In the first step, the
iridium pincer complex dehydrogenates both the PE and the light alkane
to form olefins and an Ir–H2 complex. This is followed
by the scrambling of olefins catalyzed by the olefin metathesis catalyst,
resulting in the breaking down of PE. Finally, the IrH2 complex generated in the first step transfers H2 to olefins,
resulting in the formation of alkanes. Thus, the metathesis of polyethylene
with light alkanes reduces the chain length of polyethylene and eventually
leads to the formation of shorter hydrocarbons suitable for transportation
oils. To demonstrate the proof of concept, 120 mg of high density
polyethylene [powder; MW = 3350; polydispersity index (PDI) = 1.6]
and 3 mL of n-hexane were heated at 150 °C in
a sealed vessel in the presence of the iridium catalyst Ir19 (20.1 Î¼mol), 40.2 Î¼mol of tert-butylethylene
as a hydrogen acceptor, and Re2O7/Î³-Al2O3 (57 Î¼mol of Re2O7). Analysis of the reaction mixture upon completion showed the formation
of a significant amount of C22–40n-alkanes (oil, 56% PE degradation to oils). Remarkably, no aromatic
hydrocarbon or olefins were detected by GC, although the formation
of a high molecular weight wax hydrocarbon (53 mg) insoluble in n-alkanes was obtained. Although the overall catalytic process
was successfully demonstrated, the yield of oil products was modest
(56%).</p><p>According to the predicted mechanism, a significant
reduction of
PE chain length requires the metathesis of the internal olefins generated
from PE with an internal olefin of a light alkane. Hence, the complex
(tBu2PO-tBu2POCOPtBu2)Ir(C2H4) (Ir20) was chosen as
a catalyst as it was previously reported to produce significant amounts
of internal alkenes.418 Remarkably, this
catalyst (supported on γ-Al2O3) resulted
in the formation of oil products in 98% yield. Moreover, the recyclability
of the catalyst and PE degradation at reduced catalytic loading was
also demonstrated. For practical purposes, degradation of common plastic
wastes such as postconsumer polyethylene bottles, bags, and films
was also demonstrated without any pretreatment.</p><!><p>Although
alkane metathesis is a major breakthrough in the direction of alkane
upgrading, it produces a statistical distribution of products with
limited selectivity of the desired weight fraction. A complementary
approach for alkane upgrading was reported by Bercaw and Labinger,
based on the coupling of an alkane and alkene having the same number
of carbons.419,420 The hypothesis for the tandem
catalytic system involves an alkene dimerization catalyst performing
alkene-Cn upgrading via dimerization and a transfer
hydrogenation catalyst transforming upgraded alkene-C2n to alkane-C2n. Concomitantly, an equivalent amount
of alkene-Cn will be formed during transfer hydrogenation
that will continue the catalytic cycle (Scheme 66).</p><p>For initial experiments,
Cp*TaCl2(alkene) Ta1, originally reported
by Schrock and demonstrated to be a highly
active catalyst for selective dimerization of 1-alkenes,421 and [POCOP]Ir catalyst Ir21, were
employed for the coupling of 1-hexene and n-heptane
(Scheme 66). Monitoring
the reaction by GC showed almost complete consumption of 1-hexene. n-Hexane and C12 hexene dimers were observed as the major
products. However, no C13 or C14 products were detected by GC, suggesting
that catalyst Ir21 might not be suitable for the formation
of 1-heptene. Interestingly, using a catalytic combination of Ta1 (16 mM) and the iridium pincer complex Ir22 (10 mM), known for the selective formation of terminal alkenes,
C13 and C14 alkenes were formed in a combined yield of 22%. The formation
of C12/C13/C14 alkanes was not observed. Several catalytic conditions
were screened to improve the yield of the product. A better yield
of C13/C14 alkene products was achieved by maintaining a low and steady
concentration of 1-hexene. Mechanistic investigations also suggest
that both catalysts perform independently without inhibiting each
other's catalytic activity.</p><p>In a similar direction, Goldman
and co-workers reported (in a patent)
a new strategy for alkane upgradation where an alkane is dehydrogenated
by an iridium pincer catalyst and then undergoes oligomerization producing
higher-chain olefins which subsequently are transfer hydrogenated
to form higher-chain alkanes.422</p><!><p>The use of ethylene glycol as a fuel in fuel cells has attracted
significant interest recently.423,424 In addition to potential
application in the energy sector, ethylene glycol is a valuable chemical
feedstock with applications in the manufacture of polyesters, e.g.,
PET resins (polyethylene terephthalate), and as a solvent and antifreeze
agent.423 Current industrial production
of ethylene glycol involves oxidation of ethylene to ethylene oxide
under harsh reaction conditions followed by hydrogenation. With the
constant demand for ethylene glycol, an improved and sustainable synthesis
of ethylene glycol from inexpensive and renewable feedstock is desirable.
In 2016, Beller reported a new two-step method based on oxidative
coupling of an amine with CO followed by hydrogenation to produce
ethylene glycol (Scheme 67).425 The first step was based
on an earlier report on Pd-catalyzed oxycarbonylation of amines to
oxamides by Pri-Bar and Alper.426 After
screening of several conditions, carbonylation of piperidine to 1,1-oxalyl
dipiperidine was performed under 25 bar CO and 25 bar air, catalyzed
by Pd(acac)2 (0.01 mol %), P(o-tol)3 (0.4 mol %), K2CO3 (10 mol %), and nBu4NI (2 mol %), at 120 °C
leading to a TOF up to 750 h–1 and 98% yield. The
second step was catalyzed by a combination of Ru-MACHO-BH (Ru4, 0.1 mol %) and KOtBu (2 mol %), resulting
in the complete conversion of 1,1′-oxalyl dipiperidine forming
ethylene glycol in 94% yield in 24 h (in toluene). Remarkably, an
iron-MACHO complex (2 mol %) and KOH (5 mol %) also catalyzed the
hydrogenation of 1,1′-oxalyl dipiperidine to ethylene glycol
in 77% yield. The reaction tolerated the addition of water with no
change in the catalytic activity. Moreover, separation of the desired
ethylene glycol from the reaction mixture (toluene and piperidine)
was straightforward due to the formation of a biphasic mixture, and
ethylene glycol was isolated in ∼95% yield.</p><p>Soon after, Bhanage reported a modified two-step synthesis
of ethylene
glycol using oxidative carbonylation of piperidine and ethanol followed
by subsequent hydrogenation of the product oxamates using the Milstein
catalyst Ru13 (Scheme 68).427 The first step, oxidative
carbonylation of piperidine and ethanol, was performed by an earlier
established method reported by the same group.428 Conditions for the second step, hydrogenation of oxamides,
were optimized by using several ruthenium catalysts and varying the
temperature, solvent, and base. The best results were obtained using
Milstein's RuPNN catalyst Ru13 (1 mol %) that
produced ethylene glycol in 92% yield under the conditions of 60 bar
H2, 160 °C, and 10 h in toluene solvent.</p><p>Followed by this work, Bhanage recently reported a similar two-step
synthesis of ethylene glycol with oxidative carbonylation of ethanol
using CO/O2 (25:6 atm) forming diethyl oxalate followed
by subsequent hydrogenation (Scheme 68). The first step was catalyzed by Pd/C (10 mol %)
in the presence of Bu4NI (TBAI, 0.2 mmol) at 70 °C.429 Several ruthenium catalysts were screened for
the second step: hydrogenation of diethyl oxalate to ethylene glycol
and ethanol. The best results were obtained using Milstein's
RuPNN catalyst Ru13 (1 mol %) that afforded a 92% yield
of ethylene glycol using 40 bar H2 at 100 °C in 14
h. Recyclability of the catalyst was also demonstrated for up to four
cycles without any noticeable difference in performance.</p><!><p>As discussed above, many well-defined
transition metal complexes
have been developed in recent years for their catalytic applications
in the production and storage of several energy carriers. Catalytic
(de)hydrogenative transformations to enable efficient storage of H2 have been an important focus of this review. Production of
H2 from methanol via aqueous methanol reforming reaction
has been studied well in the past using both heterogeneous and homogeneous
catalysts (section 2.1). Well-defined transition metal complexes such as ruthenium pincer
complexes represent the state-of-the-art catalysts for the low-temperature
reforming of methanol (MeOH + H2O → CO2 + 3H2, hydrogen storage capacity 12.0 wt %). As the direct
hydrogenation of CO2 to methanol has been possible industrially,330 methanol reforming/CO2 hydrogenation
is a promising approach for a reversible hydrogen storage system.</p><p>Whereas a plethora of studies has been dedicated to exploring the
dehydrogenation of aqueous methanol, dehydrogenation of formaldehyde
has received only scant attention (section 2.2). Dehydrogenation of aqueous formaldehyde
(H2CO + H2O → CO2 + 2H2, hydrogen storage capacity 8.4 wt %) is highly exothermic
(ΔHr = −35.8 kJ mol–1) compared to the dehydrogenation of aqueous methanol, which is highly
endothermic (ΔHr = 53.3 kJ mol–1), and thus provides thermodynamic advantages. However,
the reaction also poses a challenge due to catalyst poisoning from
the coordination of CO gas resulting from the decarbonylation of HCHO.
New catalyst design could facilitate the application of aqueous HCHO
for the purpose of reversible hydrogen storage.</p><p>Another potential
hydrogen storage material that has been extensively
reported in the literature and briefly discussed in this review is
amine-borane (section 2.3). Dehydrogenation of amine-boranes is thermodynamically downhill,
and a large number of homogeneous catalysts have been studied for
their dehydrogenation. However, its practical application as a hydrogen
storage material is still far away. This is because of the difficulty
in the regeneration of "charged amine-borane fuel" from
the "spent amine-borane fuel". To overcome this, cyclic
amine-boranes as a hybrid of inorganic (BN fragment) and organic (C–C
fragment) components were synthesized and dehydrogenated. However,
the hydrogenation of "spent fuel" in those cases using
molecular hydrogen has still not been demonstrated. Regeneration of
a "charged fuel" such as H3B·NH3 from direct hydrogenation of a "spent-fuel"
such as borazine is still an unsolved problem in front of the catalysis
community.</p><p>In the direction of developing reversible hydrogen
storage materials
based on organic liquids, also called liquid organic hydrogen carriers
(LOHCs), several organic compounds have been investigated (section 2.4). Dehydrogenation
of HCOOH (HCOOH → CO2 + H2) has been
well investigated; however, its reverse reaction, i.e., hydrogenation
of CO2 to HCOOH, often requires a basic medium to drive
the reaction forward by forming a formate salt. This poses a hindrance
in using HCOOH as a hydrogen storage material for commercial purposes
as the regeneration of HCOOH from a formate salt requires a stoichiometric
amount of acid. A few examples have been reported for base-free hydrogenation
of CO2 to HCOOH that uses either an acidic buffer or a
polar protic solvent such as DMSO or an ionic liquid containing an
anion that stabilizes HCOOH through H-bonding interactions. The area
of hydrogenation of CO2 to HCOOH under additive-free conditions
is still in its preliminary stage. New catalyst designs could facilitate
this transformation that could potentially make CO2/HCOOH
a commercial hydrogen storage material. Several LOHCs involving alcohols
and amines have also been developed in the past few years. Ru pincer
complexes are the state-of-the-art catalysts for such (de)hydrogenative
transformations. The advantages of using alcohols/amines as LOHCs
are that they are inexpensive and, in some cases, renewable, whose
dehydrogenation to form esters/amides, as well as the reverse hydrogenation
reactions, is facile and occurs under relatively mild conditions.
A common limitation for most of such catalytic (de)hydrogenative systems
is the use of a solvent in catalysis, thus significantly reducing
the hydrogen storage capacity. This presents opportunities to explore
organic compounds that can be dehydrogenated under neat conditions
and can be utilized as hydrogen storage materials. Renewable diols
such as ethylene glycol/polyester have also been demonstrated for
a potential LOHC (section 2.4.3). Along this line, utilization of a glycerol/polyester
couple (up to 6.52 wt %) as an LOHC can be a significant breakthrough
benefiting both the hydrogen economy and the circular economy simultaneously.
It is noteworthy that glycerol is currently being produced in a surplus
quantity as a byproduct of the biodiesel industry and therefore utilization
of glycerol to make a renewable and recyclable polyester can significantly
benefit the circular economy.430 In Table 5 we have outlined
a comparative summary of the developments and limitations of various
hydrogen carriers mediated by homogeneous catalysts discussed in this
review.</p><p>The approach of producing conventional fuels (hydrocarbons)
from
biomass such as ethanol, vegetable oil, and lignin has also been discussed
here (section 3). Several
homogeneous catalysts for upgradation of ethanol to butanol have been
reported, but selective upgradation of ethanol to higher alcohols
such as octanol and decanol has not been accomplished yet. The area
of lignin depolymerization to produce useful fuel has been mostly
studied by using heterogeneous catalysts, presenting opportunities
for new homogeneous catalyst design that could be effective for depolymerization
of lignin via C–O cleavage.</p><p>The role of homogeneous catalysts
for the methanol economy has
been discussed in detail in terms of the production of CH3OH by the hydrogenation of CO2 or CO (section 4). Direct hydrogenation of
CO2 to methanol has been mostly investigated by heterogeneous
catalysts. In fact, 100% renewable methanol is produced from the hydrogenation
of CO2 to methanol using the "emission to liquid
technology" by Carbon Recycling International. Homogenous catalysts
have been utilized for the indirect conversion of CO2 to
CH3OH, where CO2 is trapped with reagents such
as alcohol, amines, silanes or boranes, and then subsequently hydrogenated
or hydrolyzed to form methanol and regenerate the trapping agent.
The development of new homogeneous catalysts for the direct hydrogenation
of CO2 to CH3OH under mild conditions is still
a challenge in the area of molecular chemistry. An interesting development
in this direction has been made recently by Himeda and Onishi, who
have utilized the approach of a gas–solid phase reaction for
the hydrogenation of CO2 to methanol under mild conditions
(e.g., 60 °C, 40 bar, 3:1 H2/CO2, TON up
to 113).431 The reaction is catalyzed by
a multinuclear iridium complex capable of performing intramolecular
multiple hydride transfer to CO2. The success of the catalytic
methodology was partly attributed to the solid–gas phase reaction
approach that suppresses the liberation of formic acid which was found
to inhibit this transformation in the aqueous phase due to its decomposition
to H2 and CO2. Another challenging topic where
homogeneous catalysts can make a significant impact is the direct
partial oxidation of methane to CH3OH as discussed above
(section 4.4).</p><p>Alkane upgradation through the approaches of cross alkane metathesis
and alkane–alkene coupling presents attractive opportunities
to produce fuels from biomass such as CH4 (section 5). However, methane upgradation
has not been demonstrated yet due to a high barrier associated with
C–H activation of CH4, and further coupling reactions.
Moreover, the development of alkene/alkane metathesis catalysts involving
complexes of earth-abundant metals will also be a breakthrough and
enhance the sustainability of various processes based on the metathesis
reaction.</p><p>Based on the above discussions, we present here significant
challenges
in homogeneous catalysis to be achieved under relatively mild conditions,
for the development of sustainable energy carriers:</p><p>(1) direct
hydrogenation of CO2 to CH3OH
without using an additive</p><p>(2) direct hydrogenation of "spent"
B–N fuel
(e.g., aminoborane, borazine, polyaminoborane) to "charged"
B–N fuel (amine-boranes)</p><p>(3) development of glycerol/polyester
as a hydrogen storage couple
to benefit both hydrogen economy and circular economy simultaneously</p><p>(4) direct hydrogenation of CO2 to HCOOH without using
an additive or special solvent</p><p>(5) ethanol upgrading to higher
alcohols (e.g., octanol–cetyl
alcohol)</p><p>(6) C–O cleavage of aryl ethers (model lignin
compounds)
using H2 without using a stoichiometric additive</p><p>(7) methane upgradation and direct partial oxidation of methane
to methanol (CH4 + 1/2O2 → CH3OH) with high yield and selectivity</p><p>(8) alkane dehydrogenation/metathesis
reaction using earth-abundant
metal catalysts</p><p>Although a few advances have been made on some
of these topics,
we believe that further development in homogeneous catalysis can lead
to a paradigm shift in the advent of new sustainable technologies
for the production and storage of energy carriers.</p><!><p>∥ A.K. and P.D. contributed equally.</p><!><p>The authors
declare no
competing financial interest.</p><p>Amit Kumar is a Leverhulme
Trust Early Career Researcher at the
School of Chemistry, University of St. Andrews. He completed his integrated
M.Sc. chemistry degree (2007–2012) at the Indian Institute
of Technology (IIT), Roorkee, where he received several research fellowships
and awards (Indian Academy of Science, DAAD-Germany, IIT-ParisTech,
KVPY & INSPIRE from the Government of India) along with the Institute
Silver Medal. He then won a Rhodes Scholarship and pursued his DPhil
(2012–2016) under the supervision of Prof. Andrew Weller at
the University of Oxford, U.K. Upon completion of his DPhil, Amit
received the PBC fellowship (Planning & Budgeting Committee, Israel)
to carry out his postdoctoral research with Prof. David Milstein at
the Weizmann Institute of Science, Israel, where he was promoted to
be a senior postdoctoral fellow in 2019. Amit was awarded the FGS
(Feinberg Graduate School) Prize for outstanding achievements in postdoctoral
research (2018) by the Weizmann Institute of Science.</p><p>Prosenjit Daw is an assistant professor in
the Department of Chemical
Sciences at the Indian Institute of Science Education and Research
(IISER) Berhampur. He completed his B.Sc. in chemistry from the University
of Calcutta (India). He obtained his M.Sc. (2009) in chemistry and
Ph.D. (2015) from the Indian Institute of Technology Kanpur (IIT Kanpur,
India), where he worked under the supervision of Prof. Jitendra K.
Bera and received the prestigious Shyama Prasad Mukherjee Fellowship
(SPMF). After completing his Ph.D., he received the PBC fellowship
(Planning & Budgeting Committee, Israel) and moved to Israel to
work with Prof. David Milstein as a postdoctoral fellow at the Weizmann
Institute of Science (2015–2019).</p><p>David Milstein is the Israel Matz Professorial Chair
of Organic
Chemistry at the Weizmann Institute of Science. He received his Ph.D.
degree with Prof. Blum at The Hebrew University of Jerusalem in 1976
and performed postdoctoral research at Colorado State University with
Prof. Stille. In 1979 he joined DuPont Co.'s CR&D where
he became a group leader, and in 1986 he moved to the Weizmann Institute,
where he headed the Department of Organic Chemistry in 1996–2005.
In 2000, he founded and became head of the Kimmel Center for Molecular
Design until 2017. He has received several awards, including the Israel
Prize in 2012. He is a member of the Israel National Academy of Sciences
and Humanities, U.S. National Academy of Sciences, German National
Academy of Sciences—Leopoldina, European Academy of Sciences,
and is a Foreign Member of the Royal Society (ForMemRS, U.K.).</p>