This article was published as part of the

2009 Renewable Energy issueReviewing the latest developments in renewable

energy research Guest Editors Professor Daniel Nocera and Professor Dirk Guldi

Please take a look at the issue 1 table of contents to access the other reviews.

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B–N compounds for chemical hydrogen storagewzCharles W. Hamilton,

aR. Tom Baker,*

bAnne Staubitz

cand Ian Manners*

c

Received 22nd September 2008

First published as an Advance Article on the web 26th November 2008

DOI: 10.1039/b800312m

Hydrogen storage for transportation applications requires high volumetric and gravimetric

storage capacity. B–N compounds are well suited as storage materials due to their light weight

and propensity for bearing multiple protic (N–H) and hydridic (B–H) hydrogens. This critical

review briefly covers the various methods of hydrogen storage, and then concentrates on chemical

hydrogen storage using B–N compounds. The simplest B–N compound, ammonia borane

(H3NBH3), which has a potential 19.6 wt% hydrogen storage capacity, will be emphasised

(127 references).

Introduction

An abundant, inexpensive energy supply is an essential com-

ponent to boost developing economies as well as maintain

status quo for established economies. Current energy con-

sumption is based on combustion of carbon-based fuels to

water and carbon dioxide (as well as environmentally harmful

carbon particulates and sulfur and nitrogen oxides). For many

years, climate modelling has shown that increased carbon

dioxide concentration in the atmosphere will lead to detri-

mental environmental effects (recently reiterated in the fourth

IPCC report).1 Thus it is vitally important to institute a shift

away from carbon-based fuels and toward environmentally-

friendly replacements.

The transportation sector represented 31(28)% of the

EU(US) energy use in 2005(2006) and accounted for

21(34)% of CO2 emissions.2,3 For stationary energy applica-

tions, CO2 can be potentially sequestered.4 For portable

energy, however, this is impractical and an alternative energy

carrier must be used.

Hydrogen has the potential to be a clean, source-independent,

energy carrier.5 It has a high energy content per mass

compared to petroleum (120 MJ kg�1 for hydrogen versus

44 MJ kg�1 for petroleum). Also, hydrogen can be readily used

to run a fuel cell, which greatly increases efficiency over internal

combustion engines (B32% efficiency for diesel-electric; 90%

potential efficiency for fuel cell with heat capture, 85% for

electric motor, 77% overall efficiency) while eliminating for-

mation of carbon particulates and sulfur and nitrogen oxides.6

a Los Alamos National Laboratory, Inorganic, Isotope, and ActinideChemistry, MS J582, Los Alamos, NM 87545, USA.E-mail: chamilton@lanl.gov; Fax: +1 505-667-3502;Tel: +1 505-665-4636

bDepartment of Chemistry, University of Ottawa, 10 Marie Curie,Ottawa, Ontario, Canada K1N 6N5. E-mail: rbaker@uottawa.ca;Tel: +1 613-562-5800 ext. 5698; +1 613-562-5800 ext. 5613

cUniversity of Bristol, School of Chemistry, Cantock’s Close, Bristol,UK BS8 1TS. E-mail: Ian.Manners@bristol.ac.uk;Fax: +44 (0)117 929 0509w Part of the renewable energy theme issue.z Dedicated to Prof. M. Frederick Hawthorne on the occasion of his80th birthday.

Charles W. Hamilton

Charles Wayne Hamilton wasborn in Houston, Texas, USA.He received his BS in Chem-istry from Texas A&M Uni-versity in May of 2001.During his undergrad, he stu-died nuclear chemistry at theCyclotron Institute under thedirection of Prof. Joseph Na-towitz. He then joined the la-boratory of Prof. JosephSadighi at the MassachusettsInstitute of Technology wherehe studied oxidation catalysis.He received his PhD in 2007,and subsequently accepted a

post-doctoral position at Los Alamos National Laboratorywhere he currently resides. His current research interests includethe formation of molecular catalysts for amine borane dehydro-genation.

R. Tom Baker

R. Tom Baker received hisBSc degree from UBC in1975 and his PhD from UCLAin 1980 with M. FrederickHawthorne. After a postdoc-toral stint with Philip S. Skellat Penn State, he joined Du-Pont Central Research in Wil-mington, DE, where he appliedinorganic and organometallicchemistry and homogeneouscatalysis to the nylon, fluoro-products, and titanium dioxidebusinesses. He joined theChemistry Division at LosAlamos National Laboratory

in 1996 and worked on multifunctional catalysis approaches tolow-temperature hydrocarbon functionalization and chemicalhydrogen storage. In 2008, he became a chemistry professor atUniversity of Ottawa and the director of the Centre for Cata-lysis Research and Innovation.

This journal is �c The Royal Society of Chemistry 2009 Chem. Soc. Rev., 2009, 38, 279–293 | 279

CRITICAL REVIEW www.rsc.org/csr | Chemical Society Reviews

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Unfortunately, hydrogen has a poor energy content per volume

(0.01 kJ L�1 at STP and 8.4 MJ L�1 for liquid hydrogen vs.

32 MJ L�1 for petroleum).

For transportation applications, an energy carrier should

have a high energy content in as small a volume as possible to

not intrude on passenger space, and as small a mass as possible

to maintain fuel efficiency. The US Department of Energy has

established a series of targets for hydrogen storage materials

so that a vehicle can travel4500 km on a single hydrogen fill.5

Included with this are stringent system volumetric (5.4 MJ L�1

2010, 9.72 MJ L�1 2015) and gravimetric targets (6 wt% 2010,

9 wt% 2015). Evaluation of a hydrogen storage system

includes all associated components (tank, valves, regulators,

piping, mounting brackets, insulation, added coolants, etc.).

There are currently four leading methods to store hydrogen:

physical means, sorbents, metal hydrides, and so-called che-

mical hydrides. All four will be briefly summarised, and then

chemical hydrides with B–N bonds will be discussed in greater

detail. Details of the potential of ammonia borane as a

hydrogen storage material have been presented recently.7

Physical hydrogen storage

Safely containing hydrogen at high enough pressures requires

strong tanks. The tanks must also be lightweight to maintain a

high gravimetric capacity. Carbon-fibre reinforced composite

tanks are light and capable of storing hydrogen at pressures up

to 700 bar (4.7 MJ L�1). These tanks are nonconformable

(normally cylinder shaped), which makes them difficult to adapt

to a car design. There is also an energy penalty for gas

compression (15–20% of the lower heating value for hydro-

gen).5 General Motors recently demonstrated the Chevy Sequel,

which is able to achieve 482 km (300 miles) on a single hydrogen

fill using compressed tanks to store the hydrogen fuel.8 Honda

recently released the FCX Clarity, which is capable of 435 km

(270 miles) on a single fill, and has recently become available for

lease in California, USA.9 Both cars use efficient polymer

electrolyte membrane fuel cells. While these fuel cells allow

for the large range, they currently require expensive platinum

catalysts and suffer from limited durability. Although these

specific designs may not yet be practical for world-wide use,

their demonstration is an important step to show the viability of

hydrogen-fuelled cars.

While liquid hydrogen has a higher density than compressed

gas, maintaining the low temperatures requires extra compo-

nents that adversely affect system volumetric and gravimetric

storage capacity. Boil-off control is also a significant problem,

with tanks losing approximately 1% per day. Finally, signifi-

cant energy needs to be input to liquefy hydrogen (30% of the

lower heating value for hydrogen).5

Sorbents

A wide range of nanoporous materials have been studied as

potential hydrogen storage materials. The advantage of these

sorbent materials lies in their ready reversibility. This is also

their largest disadvantage, however, as they contain hydrogen

by weak physisorption forces (van der Waals), which are

normally less than 1.4 kcal mol�1. Hence, low temperatures

(normally liquid nitrogen temperature, 77 K) are necessary to

obtain reasonable hydrogen uptake.10

Carbon-based materials such as nanotubes, nanofibres,

solid foams, and activated carbon have been extensively

studied. There was initially some confusion about the hydro-

gen uptake potential of these materials due to the various

methods of measuring hydrogen uptake. Recently, the results

have been more consistent, with several methods being

employed in the same study.11 The maximum adsorption is

B5 material wt% hydrogen at 77 K; with addition of tanks

and cooling systems, the system storage capacity will be

much lower.

Zeolite structures have also been examined.12 Zeolites are

microporous structures of hydrated aluminate and silicate that

are more easily prepared than nanotubes and offer greater

control over pore size. However, they consist of atoms that are

heavier than carbon, which limits gravimetric capacity. An

emerging field is using zeolite structures as a template to form

carbon networks.13 Templating carbon onto zeolites is difficult

because zeolite structures can be disordered. Nonetheless,

Anne Staubitz obtained a diplo-

ma in biochemistry from

the University of Tubingen,

Germany. For her diploma

thesis she worked in the group

of Prof. Paul Knochel at

the University of Munich

with novel Grignard reagents

applied to the synthesis of

heterocycles. She then did her

PhD with Prof. Varinder

Aggarwal at the University of

Bristol, working on a natural

product synthesis. In Ian’s

group she is working as a

postdoctoral fellow on the catalytic dehydrogenation of group

13–group 15 adducts.

Anne Staubitz

Ian Manners received his PhD

from the University of Bristol in

1985 in the area of transition

metal chemistry. After complet-

ing postdoctoral work in Ger-

many in main group chemistry

and the USA on polymeric ma-

terials, he joined the University

of Toronto, Canada in 1990.

After 15 years he returned to

his Alma Mater to take up a

Chair in Inorganic, Macromo-

lecular and Materials Chemis-

try. His research interests focus

on the development of new

synthetic reactions in inorganic chemistry and their applications

in molecular synthesis, polymer and materials science, supramo-

lecular chemistry, and nanoscience.

Ian Manners

280 | Chem. Soc. Rev., 2009, 38, 279–293 This journal is �c The Royal Society of Chemistry 2009

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recent results demonstrated hydrogen uptake for zeolite-like

carbon materials of 6.9 wt% at 77 K and 20 bar.

Organic polymers with intrinsic microporosity (PIMs) are

another alternative.14 PIMs are made of rigid monomers to

maintain microporosity (Fig. 1a), yielding gravimetric storage

capacity of up to 3.0 wt% at 77 K and 15 bar.15 Increasing

pore size (and thus surface area) should result in a higher

storage capacity.

Metal–organic frameworks (MOFs) have also recently come

into focus as potential hydrogen storage materials.16 These are

three-dimensional polymers of metal atoms linked by bridging

ligands (Fig. 1b).17 MOFs are readily synthesised by solution

methods and then activated by heating to remove solvent

molecules. Pore sizes can be rationally designed by choosing

the metal and the linker. However, with large linkers, these

tend to form interpenetrated structures which can decrease

pore size. Similar to carbon nanotubes, MOFs rely primarily

on physisorption to bind hydrogen and low temperatures

(typically 77 K) are required.

To make these materials more practical, the temperature of

effective adsorption must be raised (ideally to ambient tem-

peratures). By adding in metal atoms or other materials

capable of chemisorption, the energy of binding can be

adjusted. This combination of chemisorption and physisorp-

tion could potentially increase the temperature of effective

uptake. There are several recent examples of increasing the

enthalpy of hydrogen adsorption for zeolites,18 MOFs,19 and

carbon materials.20 However, the inclusion of extra elements

decreases material gravimetric capacity.

Metal hydrides

Metal hydrides rely on stronger chemical interactions than

sorbents, and this results in materials that store hydrogen at

higher temperatures. Heating is required to release the hydro-

gen after uptake. Since most metal hydrides are dense pow-

ders, achieving volumetric storage targets is feasible.

Unfortunately, the relatively heavy metals make high gravi-

metric capacity difficult.

One way to achieve higher gravimetric capacity is by

including lighter main-group elements. These complex metal

hydrides, such as alanates, amide, and borohydride com-

pounds, have been evaluated as reversible hydrogen storage

materials.21 One of the early examples is titanium-doped

sodium alanate (NaAlH4), which can reversibly release hydro-

gen to give a maximum material gravimetric capacity of

5.5 wt% (although 3–4 wt% is typically obtained). Chen et al.

found that Li3N can add roughly two equivalents of hydrogen

to form lithium amide (LiNH2) and two equivalents of lithium

hydride at elevated temperatures (200–250 1C) to give a

9.3 wt% uptake of hydrogen (Scheme 1).22 Under vacuum and at

temperatures below 200 1C, 6.3 wt% of hydrogen desorbs. The

remaining 3 wt% could be removed by heating above 320 1C.

Calcium borohydride can be heated to 400 1C to release

9.6 wt% hydrogen (Scheme 1).23 Upon addition of catalytic

amounts of the dopants, TiCl3 and Pd, the ‘spent fuel’ can be

rehydrogenated at 700 bar and 400–440 1C in 60% yield.

Recent effort has concentrated on mixtures of complex

metal hydrides as potential hydrogen storage systems. Yang

and Sudik found that ternary mixtures of MgH2, LiNH2 and

LiBH4 have increased rates and extent of hydrogen release

compared to binary mixtures of these components.24 At

elevated temperatures (B350 1C), 8–11 wt% hydrogen is

released depending on the amount of MgH2 in the mixture.

At lower temperatures, hydrogen release is reversible (2.8 wt%

at 140 1C). Soloveichik and co-workers reported the decom-

position of a compound that contains borohydride and

amines, Mg(BH4)�2NH3.25 This compound has a maximum

storage capacity of 16.0 wt%. Though hydrogen loss is

observed at 150 1C, even heating to 400 1C results in a net

loss of only 13.1 wt% hydrogen. Rehydrogenation has not yet

been realised, but loss of hydrogen is reported to be endother-

mic, indicating that direct rehydrogenation may be possible.

Chemical hydrides

The so-called chemical hydrides typically use lighter elements

than metal hydrides, which result in much higher gravimetric

storage capacity. The necessary cleavage of covalent element–

hydrogen bonds, however, puts more stringent requirements

on reversibility. A reversible system will have a Gibb’s free

energy (DG) of hydrogen release at or near 0 kcal mol�1. These

materials will have a positive entropic term (DS) as hydrogengas is being released. Thus, a slightly endothermic (DH 4 0)

dehydrogenation reaction is required to achieve a system that

is reversible under reasonable/practical conditions.

Most potential chemical hydrogen storage materials suffer

from reaction enthalpies that are too endothermic or exother-

mic for reversible hydrogen release. Reaction kinetics can also

be a significant problem. Indeed, for many years, the only

known examples of hydrogen activation under mild conditions

by non-metals involved complexes only isolable in matrices.26

Scheme 1 Reactions of some complex hydrides.

Fig. 1 (a) Some examples of PIMs made up of rigid monomers.

(b) Different linkers result in different pore sizes for MOFs.17

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However, there have been several recent examples of isolable

non-metal compounds that add hydrogen under mild condi-

tions. Power and co-workers demonstrated addition of two

hydrogen molecules to a digermyne compound (Scheme 2).27

This was followed by the discovery of facile hydrogen addition

across two diaryltin molecules.27 The reverse reaction, the

conversion of two E–H bonds to hydrogen and an E–E bond,

has been demonstrated by Himmel and co-workers for a B–H

species.28 Bertrand, Schoeller, and co-workers designed

N-heterocyclic carbene analogues, with an N-aryl group sub-

stituted with a carbon group, which add hydrogen under mild

heating.29 Stephan and co-workers reported the reversible

heterolysis of hydrogen over a phosphinoborane species [1,4-

(R2P)C6F4(BR2)] to make a phosphonium-borate zwitterionic

compound.30 Finally, mixtures of Lewis acids (BR3) and Lewis

bases (PR03,31 N-heterocyclic carbenes,31,32 and NR03

33) that

are typically too sterically encumbered to form a dative bond

(frustrated Lewis pairs) are capable of adding hydrogen under

mild conditions to form borate salts. The ease of hydrogen

addition depends on the sterics and electronics of the

frustrated Lewis pairs. All of these systems have very low

gravimetric capacity, but studying the basic reactions of

hydrogen addition to non-metal systems can garner insight into

the design of reversible systems with higher storage capacity.

Cyclic organic molecules have also been evaluated as

potential hydrogen storage materials.34 Dehydrogenation of

organic molecules is typically too endothermic for hydrogen

release at moderate temperatures. However, incorporation of

fused rings and heteroatoms (O and especially N) can lower

the heat requirements for hydrogen release.35 Due to the large

kinetic barriers, a catalyst (typically noble metal-based) is also

used to aid dehydrogenation and rehydrogenation. For in-

stance, in a recent patent, Pez and co-workers reported that

N-ethylcarbazole can be hydrogenated over Ru/LiAlO2 at

1000 psi of hydrogen and 160 1C. The hydrogen can then be

released over Pd/LiAlO2 at 199 1C to give 5.6 wt% hydrogen

(Scheme 3). More importantly, the system can be cycled 5

times with no detectable degradation of the N-ethylcarbazole.

There are several systems that rely on hydrolysis of active

metals36 or chemical hydrides to produce hydrogen. Thorn

and co-workers reported that a C–H in the 2-position of

benzimidazoles is sufficiently hydridic to be protonated by a

variety of acids including water over a Pd catalyst to produce

hydrogen.37 Hydrolysis of another chemical hydride, sodium

borohydride, has also been investigated. Although it can be

stabilised in basic aqueous solutions, the addition of catalysts

or acids can be used to initiate a controlled release of hydrogen

by hydrolysis.38 Unfortunately, efficiently regenerating the

BO2� to BH4

� is difficult due to the stability of B–O bonds.

Under the current technology, chemical hydrides with high

storage capacity will require off-board regeneration, which

adds to the complexity of the hydrogen storage system.

However, these compounds could offer advantages for the

fuel distribution system. Current fuel distribution is based on

transporting liquid hydrocarbons; converting this system to

transport compressed or liquid hydrogen will be an expensive

endeavour (potentially costing trillions of dollars).39 A stable

chemical hydride could circumvent this problem, allowing for

transportation of hydrogen using the existing infrastructure.

B–N compounds

Many of the previously mentioned hydrogen storage materials

have gravimetric material capacities that may be too low for

on-board hydrogen storage requirements. However, there are

several B–N compounds with the potential to meet these

requirements. B–N compounds are well suited for hydrogen

storage because both boron and nitrogen are lightweight

elements capable of bearing multiple hydrogens. Also, B–H

and N–H bonds tend to be hydridic and protic, respectively,

resulting in normally facile hydrogen release.40 A successful

B–N hydrogen storage material must contain multiple hydro-

gen equivalents per main-group element, have a good match

between the number of hydridic B–H bonds and protic N–H

bonds, and have the stability requirements necessary for safe

storage of hydrogen. Several classes of B–N materials that

may be suitable for hydrogen storage applications will be

mentioned, and then the simplest B–N compounds, amine

boranes, will be covered in more detail.

Hydrazine borane compounds

Much of the early interest in hydrazine borane compounds

was centred around their use as propellants.41 Both hydrazine

borane and -bis(borane) have been synthesised by various

methods, the most common of which is addition of hydrazine

salts to borohydride (Scheme 4). Hydrazine borane [-bis(borane)]

has a potential hydrogen storage capacity of 13.1 [13.4] wt%

(assuming loss of 3[4] eq. of H2). Unfortunately, hydrazine

bis(borane) is a shock-sensitive explosive,42 explodes when

heated in air,43 and is therefore poorly suited for hydrogen

storage applications unless it can be stabilised.

Scheme 3 Reversible dehydrogenation of N-ethylcarbazole.

Scheme 2 Non-metal compounds that add hydrogen under mild

conditions.

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Amine triborane compounds

The first high yield synthesis of ammonia-triborane

(H3N–B3H7) was reported by Kodama and co-workers in

1959.44 Tetraborane (B4H10) is treated with tetrahydropyran

to yield tetrahydropyran-triborane and a half equivalent of

diborane. Tetrahydropyran can be readily displaced by

ammonia to yield the product. Unfortunately, tetraborane

and diborane are toxic gasses causing large-scale synthesis

by this route to be impractical. Yoon and Sneddon reported

an alternate synthesis of ammonia-triborane in which

Bu4NB3H845 salt (Scheme 5) is treated with 0.5 eq. I2 in glyme

to yield glyme–B3H7, Bu4NI and 0.5 eq. of H2.46 Similar to the

previous method, glyme is displaced by ammonia to yield the

product. The structure of ammonia-triborane was established

by X-ray methods and has a trigonal arrangement of B atoms

with one bonded to ammonia.47 Ammonia-triborane has a

storage potential of 10.6 wt% (assuming 3 eq. H2 are released),

but direct dehydrogenation methods have yet to be reported.

Kodama noted that upon treatment with sodium in ammonia,

one equivalent of hydrogen is released giving a mixture of

products including NaBH4. Sneddon found that H3N–B3H7

undergoes rapid hydrolysis after treatment with 1MHCl (7.85 eq.

of H2 after 120 min) or various metal catalysts, the best being

Rh/Al2O3 (approx. 7.5 eq. after 25 min at 21 1C with

7 mol% catalysts).46 Similar to borohydride hydrolysis, pro-

ducts with B–O bonds are produced which will be energetically

costly to regenerate.

Ammonium hydrotriborate ([NH4][B3H8]) has also been

synthesised by treatment of pentaborane with ammonium

hydroxide.48 Surprisingly, this is a stable, colourless, crystal-

line solid in pure form, and no decomposition was evident on

heating the solid at 60 1C for 70 hours. It is apparently stable

in water and alcohols and slowly decomposes to form

H3N–B3H7 and H2 when treated with benzene or ether. This

is in stark contrast to [NH4][BH4]49 which decomposes at

temperatures above �40 1C. Ammonium hydrotriborate has

a potential hydrogen storage capacity of 13.9 wt% (assuming

loss of 4 eq. H2). Hydrogen loss from alkylammonium hydro-

triborate salts has been realised by addition of metallic Si or Al

(Scheme 6).50 When [NH3Me][B3H8] is treated with Si, 5 eq.51

of hydrogen gas are released giving 9.9 wt% H2.

Guanidinium hydrotriborate, [C(NH2)3][B3H8], has been pre-

pared by treating di(guanidinium) sulfate with sodium hydro-

triborate.52 AtB100 1C, this compound violently decomposes to

form 6.2 eq. H2 (by MS), which provides 12.3 material wt%. It

was noted that some decomposition was evident if the compound

was stored at 20 1C for B2 months. Also, triaminoguanidinium

hydrotriborate, [C(NHNH2)3][B3H8] was synthesised by treat-

ment of pentaborane with triaminoguanidine.53 A decomposition

point was not reported, but a melting point of 72–78 1C was

determined. The potential hydrogen storage capacity is 11.0 wt%

(assuming loss of 8 eq. H2).

It is important to note that these polyborane compounds are

likely to be explosive, as many related compounds have been

used as rocket propellants.

Amine compounds of higher-order polyboranes

Borane has a very rich chemistry, and there are many poly-

borohydride compounds known. When ammonia is bubbled

through a solution of decaborane (B10H14), the tris(ammoniate)

of decaborane (TAD) is formed.54 This compound likely has a

[NH4][B10H13(NH3)2] formulation, and one of the NH3 groups

is only loosely bound. This formulation is thus similar to the

compound [NH2Et2][B10H13(NHEt2)].55 Although TAD is

stable under ambient conditions, upon heating to 75 1C, one

equivalent of hydrogen and one equivalent of ammonia are

released. Complete conversion to BN compounds was achiev-

able by addition of hydrazine to TAD, and heating to 800 1C

under an ammonia stream.56 It is noteworthy that decaborane

reacts vigorously with hydrazine to induce a fire in air.

In a theoretical study, Nguyen, Matus, and Dixon investi-

gated the heats of formation of several ammonium salts of

polyboranes.57 They found that formation of [B12H12][NH4]2from H2 (10 eq.), BN (2 eq.), and B (10 eq.) has a heat of

formation of only 10 to 12 kcal mol�1 (the reverse reaction

would result in loss of 8.8 wt% of H2 from [B12H12][NH4]2). In

a recent DOE progress report, Hawthorne and co-workers

investigated a series of anionic polyborane compounds for

hydrogen storage applications.58 Ammonium salts of

(B11H14)�, (B12H12)

2�, and (B10H10)2� were all synthesised

Scheme 5 Synthesis and structure of ammonia-triborane.

Scheme 4 Synthesis of hydrazine bis(borane) and hydrazine mono-

borane.

Scheme 6 Synthesis of ammonium hydrotriborate and dehydrogena-

tion of alkylammonium salts.

Scheme 7 Hydrolysis of ammonium polyborane salts.

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and hydrolyzed in the presence of a metal catalyst to yield

hydrogen, ammonium borate, and boric acid (Scheme 7).

Amine borane compounds

Amine boranes are the most thoroughly studied B–N hydro-

gen storage materials. By varying the substituents on B and N,

a variety of properties can be altered such as melting and

decomposition points as well as dehydrogenation enthalpy

and nature of the reaction products. Noth and Beyer investi-

gated the physical properties of a variety of alkylamine

boranes obtained by addition of the alkylammonium salt to

lithium borohydride (Table 1).59 An alternative way to synthe-

sise these compounds is by substitution of H3B�L (L = Me2S,

THF, Me3N) with the amine.60 The physical properties are

difficult to predict. For instance, H2EtNBH3 is one of the least

stable of the group of monoalkylamine boranes, whereas

HEt2NBH3 is one of the most stable dialkylamine boranes.

Linear H2nBuNBH3 decomposes above 10 1C, whereas the

branched isomer, H2tBuNBH3, is stable up to 120 1C.

Hawthorne reported that B-substituted amine boranes can

be prepared by reduction of alkylboroxines [(BOR)3] using

lithium aluminium hydride in the presence of trimethyl-

amine.61 The majority of the Me3NBRH2 (R = nPr, iPr, nBu,

2-Bu, iBu, tBu, n-pentyl, and n-hexyl) compounds are liquids

at room temperature with the exception of R = cyclohexyl

(mp = 40–41 1C) and benzyl (mp = 58–60 1C). Treatment of

these compounds with excess ammonia in the presence of

catalytic ammonium chloride at 100–150 1C affords B-sub-

stituted borazines, [HNBR]3, in moderate to good yield

(65–91%).62

In a theoretical study, Manners, Harvey and co-workers

investigated the effect of B- and N-substituents on the DG and

DH of dehydrogenation of HR2NBR02H.63 Loss of hydrogen

from amine boranes normally is too exergonic for reversibility

(see the section: Thermal dehydrogenation of amine boranes).

By altering the substituents (R and R0) on HR2NBR02H,

however, the DG of dehydrogenation can be made more

neutral. A strong dative N–B s-bond in the reactant

(HR2NBR02H) and a weak dative N–B p-bond in the product

(R2NBR02) results in a less exergonic dehydrogenation. In

general, the s-bond plays a more important role in determin-

ing the overall energetics, so HR2NBR02H compounds with

electron donating groups on nitrogen (resulting in a more

Lewis-basic amine) and electron withdrawing groups on boron

(resulting in a more Lewis-acidic borane) are best suited for

reversible dehydrogenation. Evaluation of a series of cyclic

amines indicates that four- and five-membered rings exhibit

very similar dehydrogenation enthalpies, whereas mechanisms

involving six-membered rings are more endothermic. This

appears to be an effect of added ring strain going from an

sp3- to an sp2-hybridised nitrogen centre.

Hydrogen loss from amine boranes is typically effected by

solvolysis (including acid- and metal-catalysed variants) or

thermolysis, in which the product distribution depends on the

reaction conditions and presence of additives or catalysts. The

easiest way to evaluate the extent of hydrogen loss is by

measuring the amount of hydrogen gas generated by either

Toepler pump or GC/MS methods. Gas burette and thermo-

gravimetric analysis (TGA) may also be used, but other

gaseous products may also contribute to the volume measured

without a method for direct identification or quantification.

Temperature-programmed desorption (TPD) provides a

means to analyze the volatile products by MS. If soluble

products are formed, NMR spectroscopy (especially 1,2H,11B, and 14,15N)64 can also give a good indication of the extent

of dehydrogenation by both the chemical shift and multiplicity

of peaks formed (Fig. 2). The wide chemical shift range for 15N

NMR is particularly useful, with sp2-hybridised N in substi-

tuted borazines from �230 to �280 ppm and sp3-hybridised N

in alkylamine boranes from �340 to �375 ppm.65 Many of

these nuclei are quadrupolar and observed line widths will

depend on the electric field gradient at the nucleus and the

nuclear correlation time. Large molecules with low symmetry

(around the quadrupolar nucleus) in viscous solvents will thus

give rise to broad resonances. The increased molecular motion

and decreased viscosity at higher temperatures can often be

used to reduce linewidths in solution NMR experiments.

Table 1 Physical properties of some alkylamine boranesa

Alkylamine borane Melting point/1C Decomp. point/1C

H3NBH3 104 B100H2MeNBH3 56 70H2EtNBH3 19 30–40H2

nPrNBH3 45 50–70H2

iPrNBH3 65 90–100H2

nBuNBH3 �48 10–15H2

tBuNBH3 96 120–140HMe2NBH3 37 150HEt2NBH3 �18 200HnPr2NBH3 30 140HiPr2NBH3 23 250HnBu2NBH3 15 120HiBu2NBH3 19 150

a Table adapted from ref. 59.

Fig. 2 Some products of dehydrogenation from AB, and 11B NMR shifts.

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Conversely, cooling the sample effectively ‘decouples’ the

quadrupolar interaction to adjacent nuclei thus reducing line

widths of, for example, 1H and 13C NMR signals in alkyl-

boranes. Finally, double rotation experiments or use of high

magnetic fields allows for useful information to be obtained

for quadrupolar nuclei in the solid state as well.66

Thermal solvolysis of amine boranes

One of the conceptually easiest ways to release hydrogen from

molecules with B–H bonds is by solvolysis (see borohydride

and polyborane hydrolysis). Reactions of amine boranes with

alcohol or water are thermodynamically downhill. However,

high temperatures (above the boiling point of water) are

necessary to induce hydrolysis under neutral or basic condi-

tions.67 In a recent article, thermally-induced solvolysis was

used in two different ways.68 The first capitalised on an

exothermic hydrogen release to induce a self-sustained reac-

tion, and the second relied on pressurising water to increase its

boiling point. A gelled mixture of AB, Al, and H2O in a 2 : 3 : 3

ratio can be ignited. The excess heat generated causes the

reaction to be self-sustaining, releasing 7.7 wt% of hydrogen

in the process. Also, a mixture of water and AB can be heated

to 135 1C (under Ar pressure) to release 3 equivalents of

hydrogen. The reaction was very clean by 11B-NMR and

MS; no borazine formation was detected.

Acid-catalysed solvolysis

Acid-catalysed hydrolysis is the oldest known process for

releasing hydrogen from amine boranes (Scheme 8).69 It is

likely that the acid functions by protonating the amine, which

releases BH3 for subsequent hydrolysis. The nature of the

amine has a profound influence on the reaction rate. For

instance, AB is hydrolysed 600 times faster than H2MeNBH3

and 4.8 � 104 times faster than HMe2NBH3. More interesting

from an engineering point of view is successful use of immo-

bilised acids such as ion exchange resins or various zeolites to

activate ammonia borane.70 Even CO2 (which generates car-

bonic acid in situ) proved to be a catalyst for the dehydrogena-

tion (albeit with a slow rate).

Metal-catalysed solvolysis

Many metals and metal complexes have been found to catalyse

amine borane solvolysis (Table 2).71–80 The recent focus in

catalyst development has been on the use of non-precious

metals. In two recent heterogeneous systems, catalyst mor-

phology greatly affected catalytic activity. Hollow spheres of

nickel metal were shown to exhibit substantial catalytic activ-

ity versus nickel powder. By doping the nickel hollow spheres

with Pt, the rate could be increased such that three equivalents

of hydrogen were released within 30 minutes.79 In a recent

study, iron nanoparticles were also found to catalyse solvo-

lysis. Borohydride reduction of Fe(SO4) results in nanoparti-

cles that slowly catalyse hydrolysis (160 min at RT with 12%

cat. loading). However, if the nanoparticles are generated in

the presence of ammonia borane, a very active catalyst forms

(8 min, RT, 12% cat. loading), which can be recycled without

loss of activity, and was similarly effective in air.80 It was

found that Fe(SO4) reduction forms crystalline material in the

absence of AB, but forms amorphous nanoparticles in the

presence of AB, which may account for the difference in

activity.

Thermal dehydrogenation of amine boranes

Although amine boranes are isoelectronic to alkanes, the

energetics and process of dehydrogenation are much different.

Dehydrogenation of ethane to ethylene, for example, is en-

dothermic (DH = 32.6 kcal mol�1) as cleavage of two strong

C–H bonds is not totally compensated for by formation of H2

and the CQC p-bond. In contrast, dehydrogenation of

ammonia borane to aminoborane (H2NBH2) is exothermic

(DH = �5.09 kcal mol�1)81 as the dative B–N bond is

converted into a stronger covalent one. Upon evaluating

the kinetics of intramolecular hydrogen loss from ammoniaScheme 8 Acid catalysed hydrolysis of ammonia borane.

Table 2 Different catalytic systems for solvolysis of amine boranes

Amine borane Catalyst Solvent Eq. of H2 released Temp/1C Time Ref.

1 H2tBuNBH3 10% Pd/C (50%wet) MeOH Approx 3 30 100 min 71, 72

Me3NBH3 20 h2 Various 10% Pd/C (50%wet) H2O, various

alcoholsHigh efficiency 20 5 min (MeOH) to 190 min

(tBuOH)73

Raney Ni (5 mol%)3 H3NBH3 Pt (20% on C) (2 mol%) H2O Approx 3 20 2 min 74

[Rh(1,5-cod)(m-Cl)]2 (2 mol%) Approx 2.7 20 minPd (2 mol%) Approx 2.5 250 min

4 H3NBH3 Dowex (12 wt%) H2O Approx 2.8 20 8 min 70CO2 No AB left 7 days

5 H3NBH3 Co (10% on C) (2 mol%) H2O Approx 2.9 20 60 min 75Ni (10% on g-Al2O3) (2 mol%) Approx 2.9 60 min

6 H3NBH3 Ni0.88Pt0.12 hollow sphere (2 mol%) H2O Approx 3 20 30 min 797 H3NBH3 Rh colloids (1 mol%) H2O Approx 2.8 20 40 s 76

Ir colloids (1 mol%) Approx 3 105 minCo colloids (1 mol%) Approx 3 60 min

8 H3NBH3 RuCl3 (0.5 mol%) MeOH Approx 3 20 5 min 779 H3NBH3 Amorphous Fe nanoparticles H2O Approx 3 20 8 min 80

10 H3NBH3 Various Co, Ni, Cu nanoparticles H2O Approx 3 20 20–300 min 78

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borane in the gas phase, Dixon and co-workers found that

the intramolecular activation barrier (32–33 kcal mol�1)82 is

actually larger than the B–N bond dissociation energy

(25.9 kcal mol�1).83 According to this result, AB should

dissociate into NH3 and BH3 before H2 loss. However, the

newly formed BH3 can catalyse AB dehydrogenation through

a six-membered transition state (Scheme 9) at a barrier only

6.1 kcal mol�1 higher in energy than separated AB and BH3.

In the condensed phase, thermolysis of amine boranes such

as AB and methylamine borane (H2MeNBH3, MeAB) have

been shown to proceed by an intermolecular mechanism that

involves initial formation of a diaminoboronium borohydride

salt (Scheme 10). Further reaction of this salt with additional

amine borane molecules builds up aminoborane chains with

formation of a new B–N bond for each hydrogen molecule

released. Computational investigations of presumed linear

polyaminoborane products, H3N(BH2NH2)nBH3, showed

that low energy coiled and helical conformations are favoured

that feature B–H� � �H–N dihydrogen bonding.84

Dehydrogenation of amine boranes typically yields a variety

of oligomeric products depending on conditions and methods

unless they are sterically blocked (Fig. 2, for AB). Dixon and

co-workers calculated the thermodynamics of the formation of

smaller oligomers [BxNxHy (x = 2, 3)] in both the gas and

condensed phase.85 Larger oligomeric products were evaluated

in the condensed phase by Miranda and Ceder.86 These

products result from both a polymeric ammonia borane cycle

(AB to PAB to PIB; see Fig. 2 for structures) and a cyclic

oligomeric pathway (AB to CTB, borazine or 1,4-polybora-

zylene). While the overall reaction enthalpies depend on

the products formed, all reactions in the study are estimated

to be mildly exothermic [�1.6 to �20 kcal mol�1 AB]. If one

considers that the entropic term contributes B8 kcal mol�1 H2,

it is clear that direct rehydrogenation will not be possible

under practical conditions and that amine boranes will need to

be regenerated in a chemical process. A few products, such as

borazine, are volatile. Loss of these products leads to con-

tamination of the hydrogen stream (potentially poisoning the

fuel cell), and material loss (limiting regeneration efficiency).

Solid state thermolysis

Initial studies of alkylamine borane (H2RNBH3) thermolysis

(at 90–120 1C) afforded mixtures of cyclic amino- and imino-

borane oligomers as well as undefined products.87 Framery

and Vaultier found that heating N- and B-substituted amine

boranes (H2RNBH3 or H3NBRH2) to 200 1C gave the corres-

ponding borazine compounds in good yield.65 Detailed studies

on the thermolysis of MeAB revealed that hydrogen is released

in two stages, one at B100 1C, and the second at 190 1C. For

the latter, a competing pathway to borazine formation was

identified as dehydrogenative cross-linking of (HMeNBH2)3 to

give an insoluble polymer.88

The parent, ammonia borane, decomposes in three distinct

steps (Scheme 11).89 The first commences at B100 1C and

peaks between 107 and 117 1C with an initial weight loss of

B1.1 eq. of dihydrogen (B7.2 wt%). The second equivalent is

lost over a much broader temperature range, with a maximum

rate at B150 1C, and the rest released at much higher

temperatures.90 Since all steps are exothermic, the high tem-

perature requirement is a reflection of the significant kinetic

barriers. The decomposition temperatures and products of

dehydrogenation are dependent on the rate that the tempera-

ture is elevated. Lower temperature ramping rates result in a

higher decomposition temperature. Analysis (IR and MS) of

the volatile thermolysis products for the first dehydrogenation

step revealed traces of B2H6, H2NQBH291 and borazine

accompanying the evolved hydrogen.92 Subsequent detailed11B solid state NMR studies at high field showed that forma-

tion of a new AB mobile phase preceded formation of the

diammoniate of diborane (DADB, [BH2(NH3)2][BH4]).

DADB, formed from two AB molecules by a hydride transfer,

actually initiates hydrogen loss and concomitant B–N bond

formation.93 A similar intermediate was proposed for the

thermal decomposition of MeAB.88 The second hydrogen-

releasing step produces cyclic iminoborane oligomers (includ-

ing borazine and B–N linked borazines, polyborazylene)

whose proposed graphitic structure is reminiscent of hexagonal

and rhombohedral phases of boron nitride.94

The dehydrogenation rate can be increased by the inclusion

of additives or by intercalation of AB in a solid scaffold.

Benedetto and co-workers found that AB samples doped or

milled with Pt (ca. 1%) had a greater extent of H2 release at

low temperatures (23% increase in H2 release at 140 1C).95

Autrey and co-workers found that a nanocomposite of meso-

porous silica and AB (1 : 1 by weight) releases hydrogen at 50 1C

with a half-reaction time of 85 min compared to a half-

reaction time of 290 min at 80 1C for neat ammonia borane.96

The peak dehydrogenation temperature was lowered from

B110 1C to B98 1C when a heating rate of 1 1C min�1 was

used. Encapsulation of AB in a 24 wt% carbon cryogel97

decreased the peak dehydrogenation temperature to B90 1C,

and there was no further decomposition at higher tempera-

tures. Volumetric measurements indicated a 9 wt% loss of

hydrogen, and no borazine formation was detected (MS).

Scheme 10 Formation of the diaminoboronium borohydride salt.

Scheme 11 Thermolysis of ammonia borane.Scheme 9 Mechanism of BH3 catalysed AB dehydrogenation.

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The mechanism(s) by which the solid scaffolds increase the

rate may be similar for both cases. It is possible that the

observed effect may be a function of increasing the surface

area of AB, as it is intercalated in nano-scale pores, which is

known to lower the phase transition temperature and thereby

presumably the dehydrogenation temperature. Another possi-

bility is that dehydrogenation of AB is catalysed by the

exposed functional groups (SiOH for silica and carboxylic

acid for carbon cryogel) on the nanocomposite. Acid catalysed

dehydrogenation of AB is well known (see below). Both

mechanisms are consistent with the observation that smaller

pore sizes reduce the decomposition temperature.

Solution thermolysis of ammonia borane

Thermal decomposition of AB in a variety of aprotic, polar

solvents is slow and results in a mixture of cyclic amino- and

iminoborane oligomeric dehydrogenation products.98 Sned-

don and co-workers found that addition of ionic liquids can

greatly increase both the rate and extent of dehydrogenation.99

While heating pure ammonia borane at 85 1C gave a negligible

amount of dihydrogen after 3 h, heating in certain ionic liquids

resulted in immediate release of dihydrogen (0.95 equivalents

after 3 h at 85 1C and even 1.5 eq. after 3 h at 95 1C, compared

to 0.8 eq. at 95 1C after 3 h for AB). Monitoring these

reactions in situ using 11B NMR provided evidence for rapid

formation and stabilisation of DADB in the ionic liquid. 11B

NMR analysis of pyridine extracts of the colorless non-volatile

residue indicated linear and branched acyclic aminoborane

chains, such as H3N(BH2NH2)nBH3 and H3NBH(NH2BH3)2,

in addition to DADB.

Acid-catalysed dehydrogenation of ammonia borane

Another way to effectively dehydrogenate AB is by the addition

of Lewis or Brønsted acids. Treatment of AB with the strong

Lewis acid B(C6F5)3 at 25 1C in ether affords the boronium

cation salt [BH2(NH3)(OEt2)][BH(C6F5)3] by hydride abstrac-

tion. Strong Brønsted acids, such as trifluoromethane sulfonic

acid (HOTf), protonate a B–H bond in AB yielding hydrogen

and the analogous boronium triflate. These boronium cations

are more reactive versions of that found in DADB and can, as a

result, initiate hydrogen release from AB even at 25 1C.

Computational studies showed that the cation interacts initially

with a B–H bond of AB, drawing a protic N–H in proximity to

a hydridic B–H, resulting in loss of hydrogen. Further mole-

cules of AB then interact similarly with the resultant cationic

complex to build the aminoborane chains stepwise (Scheme 12).

The relative concentration of acid needs to be kept low

(0.5 mol%) to avoid chain termination to aminodiborane,

B2H5(m-NH2), and concentrated solutions afford high yields

of borazine at 60 1C in 4 h.100

Anionic dehydropolymerisation of ammonia borane

In a recent DOE progress report, Sneddon and co-workers

reported that generation of catalytic amounts of H2NBH3�

increases the rate of hydrogen release.101 Metal complexes of

this anion have also been investigated as potential hydrogen

storage materials (see the Metal amidoboranes section). The

anion can be generated in situ by addition of LiNH2, LiH, or

proton sponge [1,8-bis(dimethylamino)naphthalene]. The use

of proton sponge eliminates the formation of LiBH4 and NH3

side products identified when LiNH2 or LiH was used as the

base. Although, the mechanism is currently unknown, it has

been suggested that the increased hydricity of the B–H bond in

H2NBH3� leads to facile H2 loss from its reaction with AB.102

Metal-catalysed dehydrogenation of amine boranes

Metal-catalysed dehydrogenation of amine boranes offers the

potential for additional control over both extent and rate of

hydrogen production.103–113 In a 1989 patent, Laine and Blum

claimed the dehydrogenation of several amine boranes using

Ru3(CO)12 at 60 1C (Table 3; 8–10).103 Using a heterogeneous

Pd/C catalyst, Roberts and co-workers reported conversion of

HMetBuNBH3 to the corresponding aminoborane at

120 1C.104 Despite the wide range of catalytic systems developed

subsequently (Table 3), there has yet to be a catalyst capable of

both a high rate and large extent of hydrogen release.105–113

Moreover, if these systems are to be practical, base metals

need to be used at low catalyst loadings. Finally, engineering

controlled hydrogen release has heretofore only been demon-

strated for solid catalysts so effective heterogenisation strate-

gies will be required.

Under the heavily reducing conditions of amine borane

dehydrogenation, the metal complex catalyst precursor will

often undergo changes. Frequently, the active species is much

different than the precatalyst. Manners and co-workers found

that [Rh(1,5-cod)(m-Cl)]2 catalyses the dehydrogenation of a

variety of amine boranes at room temperature or with mild

heating (Table 3; 13, 14, 18–23).106 The analogous Ir precursor

or Rh precursors with different supporting ligands exhibited

much lower activity for HMe2NBH3 (DMAB) dehydrogenation

(Table 3; 15–17, 24). The dehydrogenation of DMAB using

[Rh(1,5-cod)(m-Cl)]2 exhibits an induction period, during which

a black, opaque suspension forms. TEM analysis indicated Rh

aggregation; the UV-Vis spectrum was similar to the spectrum

of Rh colloids; addition of Hg (either at the onset, or during the

reaction progress) resulted in complete loss of activity; and

although the dark powder isolated from catalysis still had

activity, the solution after filtration had almost no activity.

These observations all point to Rh(0) colloids as the catalyti-

cally-active species. However, subsequent in situ EXAFS

(extended X-ray absorption fine structure) analysis found that

the same catalytic precursor, under different conditions, may

form soluble Rh clusters that also catalyse this reaction.107

These Rh4–6 clusters, observed in solution during the reaction,

eventually precipitate, likely due to a ligand exchange process

with the formed products. The clusters can be redissolved by

Scheme 12 A Lewis-acidic [H2BNH3]+ molecule interacts with am-

monia borane to lose hydrogen and form a new compound that is

capable of attack at two positions.

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treatment with DMAB. The process of forming colloids or

clusters can be complicated, and minor variations in conditions

can have a profound influence on metal species development.

However, it is clear in both cases that the precatalyst is

drastically altered during the reaction.

Minor alterations to supporting ligands can also greatly

influence the catalytic activity. Baker and co-workers mea-

sured the relative catalytic activity of N-heterocyclic carbene

(NHC) complexes of Ni.108 The Enders’ carbene Ni complex

was found to be 11.5 times faster than the IDipp complex and

8.8 times faster than the IMes complex.109 Also, the Enders’ Ni

complex was 4.1 times faster than a Rh(NHC) complex and

1.9 times faster than a Ru(NHC) complex. The Enders’

carbene Ni complex has the largest extent of dehydrogenation

yet seen (42.5 eq.) but requires mild heating. Initially, a minor

amount of borazine is formed. However, borazine reacts

further by crosslinking reactions to form polyborazylene.

Using borazine as the substrate under similar catalytic condi-

tions to AB results in almost no activity, indicating that AB

either activates the catalyst or is involved in the crosslinking

reactions. During the course of the reaction a dark, homo-

geneous maroon solution forms. Hg addition results in no loss

of activity, which is indicative that metal colloids are not the

active catalyst. A kinetic isotope effect (KIE) is observed for

D3NBH3, H3NBD3, and D3NBD3 (2.3, 1.7, and 3.0, respec-

tively) indicating that both N–H and B–H bond cleavage are

involved in the turnover limiting step or that N–H and B–H

bond cleavage steps have similar rates.

Solvent and substrate can also play a significant role. An

example of this was noted in the titanocene-based dehydro-

genation of amine boranes initially reported by Manners and

co-workers (Table 3; 2, 3)110 and extended by Chirik and

co-workers to include other Cp-based ligands and Zr

(Table 3; 4–6).111 Chirik and co-workers noted that

{[Z5-C5H3(SiMe3)2]2Ti}N2 has a TOF of 4420 h�1 in ben-

zene-d6 and 0.29 h�1 in THF for the dehydrogenation of

DMAB. Also, AB was found to have a much slower rate than

DMAB (Table 3; 4 and 5).

Finally, the catalyst identity controls the extent of dehy-

drogenation. Heinekey, Goldberg and co-workers reported

that (POCOP)IrH2112 is an extremely active catalyst for AB

dehydrogenation (Table 3; 25).113 Unfortunately, only a single

equivalent of hydrogen is released per equivalent of AB.

During the course of the reaction, an insoluble colorless

precipitate is formed. The X-ray powder diffraction and IR

data agree closely with that previously reported for cyclopen-

taborazane, [H2NBH2]5.114 However, Manners et al. found

that alkylamine boranes, and mixtures of alkylamine boranes

and AB under similar conditions gave soluble aminoborane

polymers.115 The measured wide angle X-ray scattering pat-

tern and the IR for the white precipitate from the (POCOP)IrH2

catalysed dehydrogenation of AB were different from that

reported for [H2NBH2]5. They suggested that the white

precipitate may be polymeric in nature similar to the products

formed from alkylamine borane dehydrogenation. Although

the actual nature of the precipitate has not yet been completely

Table 3 Selection of reported metal catalysts for amine borane dehydrogenation

Catalyst (mol%) Substrate Conditions Products Eq. of H2 Ref.

1 Cp2TiMe2 (0.5%) HMe2NBH3 16 h, 25 1C No reaction 0 1062 Cp2Ti (2%) HMe2NBH3 4 h, 20 1C (Me2NBH2)2 1 1103 Cp2Ti (2%) H(i-Pr)2NBH3 1 h, 20 1C iPr2NBH2 1 1104 {[Cp(SiMe3)2]2Ti}N2 (2%) HMe2NBH3 7 min, 23 1Ca (Me2NBH2)2 1 1115 {[Cp(SiMe3)2]2Ti}N2 (2%) H3NBH3 161 h, 65 1Ca CTBb, borazine NRc 1116 [indenyl-(SiMe3)2]2Zr HMe2NBH3 147 h, 65 1Ca (Me2NBH2)2 1 1117 [P(iPr)3]2Br2(CH3CN)(NO)Re (1%) HMe2NBH3 4 h, 85 1C (Me2NBH2)2 1 1058 Ru3(CO)12 (0.2%) H3NBH3 85 h, 60 1C BN1.13H4.7 (elemental analysis) NR 1039 Ru3(CO)12 (0.1%) Me3NBH3, PrNH2 32 h, 60 1C [–N(Pr)B(H)–]3 (57%) NR 103

10 Ru3(CO)12 (0.1%) Me3NBH3, MeNH2 9.5 h, 60 1C [–B(NMeH)N(Me)–]3, B(NHMe)3d NR 103

11 trans-RuMe2(PMe3)4 (0.5%) HMe2NBH3 16 h, 25 1C (Me2NBH2)2 1 10612 FeH(PMe2CH2)(PMe3)3 (9%) H3NBH3 96 h, 25 1C CTB, borazine, polyborazylened NR 10813 [Rh(1,5-cod)(m-Cl)]2 (0.5%) HMe2NBH3 8 h, 25 1C (Me2NBH2)2 1 10614 [Rh(1,5-cod)(m-Cl)]2 (5%) HMe2NBH3 o2 h, 25 1C (Me2NBH2)2 1 10615 RhCl3 (0.5%) HMe2NBH3 22.5 h, 25 1C (Me2NBH2)2 (90%) 0.9 10616 HRh(CO)(PPh3)3 (0.5%) HMe2NBH3 160 h, 25 1C (Me2NBH2)2 (5%) 0.05 10617 [Cp*Rh(m-Cl)Cl]2 (0.5%) HMe2NBH3 112 h, 25 1C (Me2NBH2)2 1 10618 [Rh(1,5-cod)(m-Cl)]2 (0.5%) H(1,4-C4H8)NBH3 24 h, 25 1C [(1,4-C4H8)NBH2]2 (73%)e NR 10619 [Rh(1,5-cod)(m-Cl)]2 (0.5%) HMe(PhCH2)NBH3 24 h, 25 1C [Me(PhCH2)NBH2]2 (79%)e NR 10620 [Rh(1,5-cod)(m-Cl)]2 (1%) H2MeNBH3 B60 h, 45 1C (MeNBH)3 (40%)e NR 10621 [Rh(1,5-cod)(m-Cl)]2 (0.6%) H2PhNBH3 16 h, 45 1C (PhNBH)3 (56%)e NR 10622 [Rh(1,5-cod)(m-Cl)]2 (0.6%) H3NBH3 B60 h, 45 1C Borazine (10%),e PIB, polyborazylene NR 10623 [Rh(1,5-cod)(m-Cl)]2 (1%) HiPr2NBH3 24 h, 25 1C (iPr)2NBH2 (49%)e NR 10624 [Ir(1,5-cod)(m-Cl)]2 (0.5%) HMe2NBH3 136 h, 25 1C (Me2NBH2)2 (95%) 0.95 10625 (POCOPf)Ir(H)2 (0.5%) H3NBH3 14 min, 20 1C Cyclopentaborazane 1 11326 Ni(1,5-cod)2, 2 NHCg (9%) H3NBH3 3 h, 60 1C Polyborazylened 2.8 10827 Pd/C HMetBuNBH3 1 h, 120 1C [MetBuNBH2]2 1 10428 Pd/C (0.5%) HMe2NBH3 68 h, 25 1C (Me2NBH2)2 (95%) 0.95 10629 (IDipp)hCuCl (12.5%) HMe2NBH3 24 h, 20 1C (Me2NBH2)2

d NR 108

a Inferred from reported TOF values. b CTB is cyclotriborazane. c NR is ‘‘not reported.’’ d Yield of products not quantified; other products

possible. e Isolated yields, actual yield will be higher but other products detected. f k3-2,6-[OP(t-Bu)2]2C6H3.g Enders’ carbene: (1,3,4-triphenyl-

4,5-dihydro-1H-1,2,4-triazol-5-ylidene). h IDipp is 1,3-bis(2,6-diisopropylphenyl)-1,3-dihydro-2H-imidazol-2-ylidene.

288 | Chem. Soc. Rev., 2009, 38, 279–293 This journal is �c The Royal Society of Chemistry 2009

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established, it is clear from the solid state 11B NMR data

(�18 ppm) and the measured amount of hydrogen released

that an [H2NBH2]n product is formed.

The mechanisms of metal-catalysed dehydrogenation have

been investigated by computational methods for the titanocene

system reported by Manners, the (POCOP)IrH2 system

reported by Goldberg and Heinekey, and the Ni(NHC)2system reported by Baker. Although the relative energies of

intermediates are different, the overall reaction pathways are

similar. In the initial step, B–H coordinates to the metal

complex. In the titanocene case, the N–H is activated followed

by hydride transfer from B–H (Scheme 13).116 Two mechan-

isms were evaluated in the (POCOP)IrH2 case. The first

follows generation of monovalent (POCOP)Ir with a con-

certed B–H and N–H bond activation at the IrI centre. The

second one involves a concerted B–H and N–H bond activa-

tion at trivalent (POCOP)IrH2 (Scheme 14). In the second

mechanism, a hydride ligand acts as a proton acceptor yielding

pentavalent (POCOP)IrH4.117 In the Ni(NHC)2 case, N–H

deprotonation is proposed to occur at the Ni-bound NHC

carbon, forming a coordinated imidazolium-type ligand

(Scheme 15).118 The imidazolium subsequently protonates

the Ni centre and hydrogen formation follows. The proposed

mechanism is not consistent with the observed KIE, so further

investigation is warranted.

All mechanisms proposed to date are focused on the initial

loss of hydrogen from amine boranes. Establishing the role of

the metal complex in subsequent oligomerisation and loss of

the second equivalent of hydrogen will also be very important.

Regeneration of spent fuel

Spent fuel from solvolysis methods

The advantage of solvolysis methods is their high efficiency,

chemical robustness, and the fact that inexpensive base-metal

catalysts have been developed. A significant drawback of these

systems is that B–H bonds are converted to much stronger

B–O bonds, resulting in a more exothermic reaction than

dehydrogenation. Regeneration of spent fuel from rehydro-

genation methods requires strong reducing agents. Ramachan-

dran and co-workers demonstrated a system based on transition

metal catalysed solvolysis of AB to yield [NH4][B(OMe)4],

which could be converted back to AB by treatment with NH4Cl

and lithium aluminium hydride (Scheme 16).77 It will likely be

energetically costly to convert the oxidation product,

Al(OMe)3, back to the complex hydride.

Spent fuel from dehydrogenation methods

The strategies presented to date for regeneration of BNHx-

spent fuel involve two important steps, digestion and

Scheme 13 Mechanism of TiCp2 catalysed dehydrogenation of

Me2HNBH3 as proposed by Luo and Ohno.116 Gibbs free energies

corrected for toluene (CPCM).

Scheme 14 Relative energies of the proposed catalytic cycle in THF

(CPCM).

Scheme 15 Simplified partial mechanism for NHC2Ni catalysed

dehydrogenation of AB. (i) Protonation of an NHC ligand.

(ii) Transfer of a proton to metal centre. (iii) Formation of hydrogen.

Scheme 16 Solvolysis and subsequent regeneration of ammonia

borane.

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reduction. Although a multitude of products can result from

AB dehydrogenation, they all contain newly formed B–N

linkages that must be broken. Addition of an acid (HX) can

protonate these linkages, releasing amine and making B–X

bonds (digestion). The B–X bonds can then be reduced by

chemical reductants to form B–H bonds. A practical system

must be as energetically efficient as possible, so no reaction

step can be too exo- or endothermic.

The method mentioned above by Ramachandran and co-

workers could potentially be used to regenerate ammonia

borane. However, the generation of strong B–O bonds in the

reaction pathway will likely limit overall efficiency. Alterna-

tively, Sneddon and co-workers102 andMertens and co-workers119

have independently developed reduction schemes that form

B–Br and B–Cl bonds, respectively.

Sneddon and co-workers reported that spent BNHx fuel was

successfully digested using HBr/AlBr3 (super-acid) in CS2 to

form H2, NH4Br, BBr3 as well as [H2NBBr2]3. The relative

ratio of BBr3 to [H2NBBr2]3 depends on the material being

digested. Mertens and co-workers treat a THF solution of

spent fuel with an ether solution of HCl to generate BCl3,

NH4Cl and H2. Both aminoborane [(BH2NH2)x] and imino-

borane [(BHNH)x] materials were successfully digested. Un-

fortunately, the yields of BCl3 were low due to decomposition

in THF. Switching to a similar super-acid solution that

Sneddon and co-workers used increases the yield of BCl3 to

460%. BCl3 is difficult to directly hydrodechlorinate and

reduction by hydrogen requires high temperatures (600–700 1C)

to yield the partially reduced product, BHCl2 (which must be

removed from the side-product HCl and can subsequently

disproportionate into BCl3 and B2H6). However, addition of a

Lewis-base such as NMe3 reduces the hydrodechlorination

temperature to 200 1C (but still requires high pressures of

2000 atm).120 Unfortunately, these conditions are energetically

costly, and the yield is poor (25%). However, a similar concept

can be used for reduction of BX3 compounds by chemical

hydrides. Another advantage offered by a Lewis-base is that it

eliminates formation of B2H6, a hazardous material. Thus,

Sneddon and co-workers treat BBr3 with N,N0-diethylaniline,

which yields a compound easily reduced by triethylsilane

under mild conditions. Mertens and co-workers found that

stronger bases such as NEt3, inhibit the reduction of BCl3 by

triethylsilane and MgH2. Addition of a weaker base, NPh3,

allows for the complete reduction of Ph3NBCl3 by MgH2 after

20 minutes at 80 1C. The final step is displacement of R3N with

ammonia, to yield H3NBH3.

In a DOE progress report regarding Los Alamos National

Laboratory’s effort toward BNHx-spent fuel processing, effec-

tive digestion of polyborazylene with 1,2-benzenedithiol was

demonstrated to form ammonia adducts of dithioboron

compounds.121 These compounds contain relatively weak

B–S bonds that are readily reduced by Sn–H. Optimisation

of reaction conditions and energy efficiency are still being

developed, but this is an intriguing result.

Horizons

Despite the vast amount of progress made on understanding

and controlling the properties of B–N compounds, more work

needs to be accomplished for these materials to be practical

hydrogen storage materials. From an engineering point of

view, fuels and spent fuels that are liquids are preferable,

especially for materials that require off-board regeneration.

A tailored mixture of B–N compounds, such as substituted

amine boranes, could possibly be mixed with AB to afford

such a liquid fuel that could still maintain acceptable gravi-

metric storage capacity. Combinations of this fuel mix with

solid catalysts that maintain acceptable extent and rate of

hydrogen release would still need to generate a pure hydrogen

stream for the fuel cell and a spent fuel that can be regenerated

in high yield with minimum energy input. Given the energy

required just to transport the spent fuel for chemical proces-

sing, research is underway to discover new B–N hybrid

compounds that are capable of on-board reversible hydrogen

storage. Two promising examples are metal amidoboranes and

C–B–N compounds.

Metal amidoboranes

Just as complex hydrides use main-group elements to increase

the gravimetric capacity of metal hydrides, metal amido-

boranes use metals to potentially decrease the enthalpy of

hydrogen loss, eliminate formation of volatile coproducts

(increasing purity of hydrogen stream), and help control the

rate of H2 release. AB is readily deprotonated by a variety of

metal hydrides to form a new family of metal amidoboranes.

MI(NH2BH3) (MI = Li, Na) were prepared by ball-milling

solid AB with sodium- or lithium hydride.122 Thermal decom-

position of these materials was monitored using TPD, coupled

with an MS analyser. Under the conditions of measurement,

the Li and Na compounds released hydrogen at considerably

lower temperatures (89 and 92 1C, respectively) than AB

(108 1C). Formation of borazine, observed at 154 1C for AB,

was not observed in the case of the alkali-metal amidoboranes.

Both compounds have shorter B–N bond distances than that

in AB, which may be an indication of a stronger dative

bond.123 The different dative bond may be responsible for

the different decomposition temperature, but this is still a

matter of speculation.

Burrell and co-workers synthesised Ca(NH2BH3)2 by addi-

tion of AB to CaH2 in THF.124 A THF adduct is formed, but

the THF can be removed in vacuo. While this compound is

reported to lose hydrogen at 90 1C, only o0.3 equivalents of

H2 are released. Slowly ramping the temperature to 170 1C

results in 3.6 equivalents of H2 released (corresponding to

7.2 material wt%) with o0.1% ammonia and borazine released.

In addition to expanding the range of metals used in metal

amidoborane complexes, the role of dopants or catalysts in

controlling dehydrogenation, and possibly rehydrogenation,

needs to be explored.

C–B–N compounds

Another strategy to obtain reversible chemical hydrogen

storage materials is through C–B–N compounds. Since dehy-

drogenation of C–C bonds is significantly endothermic, new

compounds containing both C–C and B–N bonds may be

tailored to undergo dehydrogenation with DG E 0. Indeed,

computational results from Dixon and co-workers show that

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replacement of 2 methylene groups of cyclohexane with an

H2NBH2 group gives a material with 7.1 wt% hydrogen for

loss of 3 eq. of H2 and DH = 23.5 kcal mol�1 H2

(Scheme 17).125 Similar compounds have been reported

recently and identification of suitable dehydrogenation cata-

lysts is in progress.126

Conclusions

The high gravimetric capacity of B–N materials makes them

particularly appealing for hydrogen storage applications. It is

important to note that basic research in this area also impacts

several other fields such as transfer hydrogenation,33,127 cera-

mic precursor production, and B–N polymer synthesis.115

A wide range of materials have been synthesised, and our

understanding of hydrogen release mechanisms is rapidly

increasing. As the next generation of hybrid B–N materials

are designed to enable energy efficient regeneration or even

reversible storage, these compounds continue to show promise

for practical hydrogen storage.

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90 A good indication of whether pure boron nitride has beenobtained is the IR spectrum: often, authors cite WAXS data asproof, where BN can indeed be identified, but some N–H bondsremain, which can be identified by a prominent N–H stretch bandaround 3600 cm�1.

91 Both compounds were confirmed as volatile products in matrixisolation studies: J. D. Carpenter and B. S. Ault, Chem. Phys.Lett., 1992, 197, 171.

92 P. M. Kuznesof, D. F. Shriver and F. E. Stafford, J. Am. Chem.Soc., 1968, 90, 2557.

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94 The thermal dehydrogenation of borazine to form boron nitridehas elicited considerable interest due to the material properties ofboron nitride. While this expands the chemical knowledge offormally the third dehydrogenation step of ammonia borane, thisis of no interest for hydrogen storage purposes, as firstly, borazineis a fuel cell poison and secondly, formation of boron nitrideas a thermodynamically extremely stable product makesrecycling difficult. Readers who are interested in this aspect ofammonia borane chemistry are referred to, for example:A. W. Laubengayer, P. C. Moews, Jr and R. F. Porter, J. Am.Chem. Soc., 1961, 83, 1337; P. J. Fazen, J. S. Beck, A. T. Lynch,E. E. Remsen and L. G. Sneddon, Chem. Mater., 1990, 2, 96;P. J. Fazen, E. E. Remsen, J. S. Beck, P. J. Carroll, A. R. McGhieand L. G. Sneddon, Chem. Mater., 1995, 7, 1942; D.-P. Kim,K.-T. Moon, J.-G. Kho, J. Economy, C. Gervais andF. Babonneau, Polym. Adv. Technol., 1999, 10, 702.

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98 J. S. Wang and R. A. Geanangel, Inorg. Chim. Acta, 1988, 148,185; W. J. Shaw, J. C. Linehan, N. K. Szymczak,D. J. Heldenbrandt, C. Yonker, D. M. Camaioni, R. T. Bakerand T. Autrey, Angew. Chem., Int. Ed., 2008, 47, 7493.

99 M. E. Bluhm, M. G. Bradley, R. Butterick III, U. Kusari andL. G. Sneddon, J. Am. Chem. Soc., 2006, 128, 7748.

100 F. H. Stephens, R. T. Baker, M. Hernandez-Matus, D. J. Grantand D. A. Dixon, Prepr. Pap. - Am. Chem. Soc., Div. Fuel Chem.,2006, 51, 573.

292 | Chem. Soc. Rev., 2009, 38, 279–293 This journal is �c The Royal Society of Chemistry 2009

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101 L. G. Sneddon, Amineborane Hydrogen Storage - New Methodsfor Promoting Amineborane Dehydrogenation/Regeneration Reac-tions, DoE Hydrogen Annual Progress Report, 2007 (http://www.hydrogen.energy.gov/pdfs/progress07/iv_b_5e_sneddon.pdf).

102 L. G. Sneddon, Amineborane Based Chemical Hydrogen Storage,DoE Hydrogen Annual Merit Review, 2007 (http://www.hydrogen.energy.gov/pdfs/review07/st_27_sneddon.pdf).

103 Y. D. Blum and R. M. Laine, US Pat., 4 801 439, 1989.104 I. G. Green, K. M. Johnson and B. P. Roberts, J. Chem. Soc.,

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107 Y. Chen, J. L. Fulton, J. C. Linehan and T. Autrey, J. Am. Chem.Soc., 2005, 127, 3254; J. L. Fulton, J. C. Linehan, T. Autrey,M. Balasubramanian, Y. Chen and N. K. Szymczak, J. Am.Chem. Soc., 2007, 129, 11936.

108 R. J. Keaton, J. M. Blacquiere and R. T. Baker, J. Am. Chem.Soc., 2007, 129, 1844.

109 Enders’ carbene: 1,3,4-triphenyl-4,5-dihydro-1H-1,2,4-triazol-5-ylidene; IMes: 1,3-bis(2,4,6-trimethylphenyl)-1,3-dihydro-2H-imidazol-2-ylidene; IDipp: 1,3-bis(2,6-diisopropylphenyl)-1,3-di-hydro-2H-imidazol-2-ylidene.

110 T. J. Clark, C. A. Russell and I. Manners, J. Am. Chem. Soc.,2006, 128, 9582.

111 D. Pun, E. Lobkovsky and P. J. Chirik, Chem. Commun., 2007,3297.

112 POCOP is k3-2,6-[OP(t-Bu)2]2C6H3.113 M. C. Denney, V. Pons, T. J. Hebden, D. M. Heinekey and

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K. Diyabalanage, J. C. Gordon, C. W. Hamilton, M. Inbody,K. K. Jonietz, R. J. Keaton, V. Pons, T. A. Semelsberger,R. Shrestha, F. H. Stephens, D. L. Thorn and W. Tumas,Chemical Hydrogen Storage Research at Los Alamos NationalLaboratory, DoE Hydrogen Annual Progress Report, 2007 (http://www.hydrogen.energy.gov/pdfs/progress07/iv_b_5g_ott.pdf).

122 Z. Xiong, C. K. Yong, G. Wu, P. Chen, W. Shaw, A. Karkamkar,T. Autrey, M. O. Jones, S. R. Johnson, P. P. Edwards and W. I.F. David, Nat. Mater., 2008, 7, 138.

123 A shorter bond distance on its own cannot per se serve as prooffor a stronger bond. For a discussion see for example: F. Bessacand G. Frenking, Inorg. Chem., 2006, 45, 6956.

124 H. V. K. Diyabalanage, R. P. Shrestha, T. A. Semelsberger,B. L. Scott, M. E. Bowden, B. L. Davis and A. K. Burrell, Angew.Chem., Int. Ed., 2007, 46, 8995.

125 A. J. Arduengo and D. A. Dixon,Main Group Element and OrganicChemistry for Hydrogen Storage and Activation, DoE HydrogenProgram Review, 2008 (http://www.hydrogen.energy.gov/pdfs/review08/st_9_dixon.pdf); K. Goldberg and M. Heinekey, Solutionsfor Chemical Hydrogen Storage: Dehydrogenation of B–N Bonds,DoE Hydrogen Program Review, 2008 (http://www.hydrogen.energy.gov/pdfs/review08/st_10_goldberg.pdf).

126 E. R. Abbey, L. N. Zakharov and S.-Y. Liu, J. Am. Chem. Soc.,2008, 130, 7250; E. R. Abbey, J. T. Jenkins, L. N. Zakharov andS.-Y. Liu,Org. Lett., 2007, 9, 4905; M. Scheideman, G. Wang andE. Vedejs, J. Am. Chem. Soc., 2008, 130, 8669.

127 C. A. Jaska and I. Manners, J. Am. Chem. Soc., 2004, 126, 2698;D. Chen and J. Klankermayer, Chem. Commun., 2008, 2130.

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