FEATURE ARTICLE
www.rsc.org/materials | Journal of Materials Chemistry
Hydrogen storage in metal–organic frameworks{
David J. Collins and Hong-Cai Zhou*
Received 26th February 2007, Accepted 21st March 2007
First published as an Advance Article on the web 11th April 2007
DOI: 10.1039/b702858j
For any potential hydrogen-storage system, raw uptake capacity must be balanced with the
kinetics and thermodynamics of uptake and release. Metal–organic frameworks (MOFs) provide
unique systems with large overall pore volumes and surface areas, adjustable pore sizes, and
tunable framework–adsorbate interaction by ligand functionalization and metal choice. These
remarkable materials can potentially fill the niche between other physisorbents such as activated
carbon, which have similar uptake at low temperatures but low affinity for hydrogen at ambient
temperature, and chemical sorbents such as hydrides, which have high hydrogen uptakes but
undesirable release kinetics and thermodynamics.
Introduction
Interest is high worldwide in replacing the burning of fossil
fuels with generation of energy through hydrogen fuel cells.
Electricity generation from hydrogen and oxygen, especially
for transportation applications, has the potential to result in a
significant reduction of CO2 emissions and air pollution in
general. Of course, any fuel, especially one for use in transportation applications, would be useless without an effective
storage technology. The United States Department of Energy
has set a number of short-term goals for on-board hydrogen
storage systems: 6.0 wt% and 45 g L21 by the year 2010, and
9.0 wt% and 81 g L21 by the 2015.1 Additionally, these goals
should be met at near-ambient temperatures and applicable
pressures (less than 100 bar), and any storage technology must
Department of Chemistry and Biochemistry, Miami University, Oxford,
OH, USA. E-mail: zhouh@muohio.edu; Fax: +1 513 529-0452;
Tel: +1 513 529-8091
{ This paper is part of a Journal of Materials Chemistry theme issue on
New Energy Materials. Guest editor: M. Saiful Islam.
David J. Collins
David Collins received his B.S.
(Materials Science & Engineering) from Ohio State University
in 1995. After graduation, he
taught high school physics and
chemistry in Houston, Texas,
and Cincinnati, Ohio, before
returning to school to pursue a
Ph.D. in chemistry in 2003. At
Miami University, working with
Professor Hong-Cai Zhou,
David has investigated the
synthesis of heterometallic
paddlewheel clusters and the
synthesis and properties of novel
MOF materials, while also
pursuing a project to bring
inorganic chemistry and crystallography concepts to local high
school classes.
3154 | J. Mater. Chem., 2007, 17, 3154–3160
minimize weight and volume contributions to the potential
fuel-cell powered vehicle. To maintain current transportation
ranges, the average light-duty vehicle would require 5–15 kg of
onboard hydrogen storage.2 While liquid fuels such as gasoline
and diesel fuel can be easily stored in simple tanks at ambient
pressures, a gaseous fuel such as hydrogen poses a real
challenge. One alternative is liquefaction; however, this
requires cryogenic temperatures, extremely high pressures, or
both; compressed H2 technologies have addressed storage
pressures of 700 bar at room temperature. In addition, the
density of liquid H2 is 70.8 g L21; the 2015 gravimetric goal
(81 g L21) precludes any purely physical storage of molecular
dihydrogen as a liquid or compressed gas. Aside from
improvements in current tank technology, the storage methods
presently under the most consideration are physisorbents and
chemical and metal hydrides.
In addition to total storage capacity, the kinetics and
thermodynamics of release and recharging must be considered.
Undesirable kinetics and thermodynamics can not only limit
the release and recharging rate, but may also add unnecessary
Hong-Cai Zhou
Hong-Cai
‘‘Joe’’
Zhou
obtained his B.S. from Beijing
Normal University, and his
Ph.D. from Texas A&M
University, under the direction
of Professor F. A. Cotton,
followed by a postdoctoral stint
in the laboratory of Professor
R. H. Holm at Harvard
University. He is currently
an associate professor in the
Department of Chemistry
and Biochemistry at Miami
University. His research is in
the fields of biomimetic materials chemistry and bio-inspired
coordination chemistry.
This journal is ß The Royal Society of Chemistry 2007
weight and volume to the vehicle because of the need for extra
heat exchangers—this leads to consideration of the heat of
formation of the hydrogen-carrying species or heat of
hydrogen adsorption of sorbents, DHf or DHads, when
evaluating a potential storage method. Chemisorption techniques, involving the formation of hydrides, suffer from binding
hydrogen too tightly, i.e., DHf too large—ranging from 50 to
over 200 kJ mol21. For metal hydrides, such as LaNi5H6.5, the
storage system must operate above ambient temperature (often
.400 K) in order to release the hydrogen fuel, while refueling
liberates large amounts of energy.2 The thermodynamics of
chemical hydrides (such as LiBH4) are such that there are
questions regarding the reversibility of the hydrogen uptake
and release, while the chemical reaction which releases
hydrogen only occurs at elevated temperatures; additionally,
the energy input to create the hydride reduces the overall
efficiency of the system.2 As a contrast, most physisorption
techniques, including carbon nanotubes and other porous
materials (activated carbon, zeolites, etc.), suffer from the
opposite problem: DHads is typically small, less than
10 kJ mol21, and appreciable adsorption can only be achieved
at very low temperatures (typically ,100 K).
MOFs as physisorbent materials
In the last 10 years, metal–organic frameworks (MOFs, also
known as coordination networks or coordination polymers)
have become a burgeoning field of research and a promising
candidate for hydrogen storage materials due to their
exceptionally high porosity, uniform but tunable pore size
and well-defined hydrogen occupation sites. Generally speaking, these materials are constructed by coordinate bonds
between multidentate ligands and metal atoms or small metalcontaining clusters (referred to as the secondary building unit
or SBU, a term borrowed from the description of zeolites).
Most have three-dimensional structures incorporating uniform
pores and a network of channels. These pores and channels are
often filled with guest species, usually solvent from synthesis.
Removal of these guests often leads to framework collapse;
however, in many cases, framework integrity is preserved, and
these voids remain and other guest molecules can then be
adsorbed onto this porous structure.3
In May 2003, Yaghi, et al. reported what is believed to be
the first measurements of hydrogen adsorption on a MOF: a
remarkable 4.5 wt% at 77 K and pressures less than 1 atm, and
1.0 wt% at room temperature at 20 bar on the material
Zn4O(bdc) (bdc = 1,4-benzenedicarboxylate) (also referred to
as MOF-5 and IRMOF-1).43 These values were later adjusted
downward based on follow-up studies,12 but the idea
remained: for hydrogen storage, these porous MOFs were a
competitive alternative to other physisorption-based materials
such as zeolites or activated carbon. Since 2003, at least 60
unique MOFs have been evaluated for their ability to store
hydrogen (summarized in Table 1). Coupled with measurements of porosity and surface area based on nitrogen adsorption, some understanding of the many factors that determine
the hydrogen uptake by a porous MOF has been developed.
Recent computational studies, involving both electronicstructure methods (ab initio and DFT) and molecular
This journal is ß The Royal Society of Chemistry 2007
mechanics (Grand Canonical Monte Carlo methods), in
addition to increasingly detailed structural characterization
of the hydrogen-adsorbed species (including neutron scattering
and synchrotron X-ray diffraction) have added insight to
these remarkable materials and the mechanisms of hydrogen
adsorption.
Rather than present merely a summary of all studies that
have been performed to date, this Feature Article will focus on
those characteristics of MOFs which must be achieved and/or
improved in order to reach the promise of hydrogen storage at
near-ambient temperatures and reasonable pressures. These
factors include achieving sufficient surface area and pore
volume, the inclusion of appropriately sized pores, and the
formation of high-energy hydrogen binding sites (on either the
metal cluster, ligand, or both).
The first factors, surface area and pore volume, scale with
the overall hydrogen saturation uptake; existing studies
(vide infra) show that the surface areas and porosity achieved
currently by MOFs can nearly reach the 2010 US DOE
gravimetric adsorption goals, albeit at 77 K. The remaining
factors influence the interaction between the framework and
the hydrogen molecule, DHads; their effects are observable
primarily at low pressures and temperatures. In this regime,
the pores remain mostly unfilled; hydrogen–framework interactions dominate the shape of the adsorption isotherm.
The most facile method to determine the heat of adsorption,
DHads, is application of the Clausius–Clapeyron equation to
adsorption data collected at two temperatures (typically 77 K
and 87 K).44 Other methods include microcalorimetry and
estimation via simulation (GCMC, etc.). As can be seen in
Table 1, only about one-third of published hydrogen adsorption studies include DHads data; increasing this fraction is
necessary to further elucidate the relationships between
various structural and chemical factors and this key variable.
Since many studies do not directly measure DHads, the
hydrogen uptake at 77 K and 1 atm can be used as a
comparative proxy for the heat of adsorption. Maximizing this
hydrogen affinity will increase the temperatures at which
MOFs can adsorb large amounts of hydrogen; this is necessary
in order to develop a storage system that meets the above
targets at near-ambient temperatures.
Pore volume and surface area
In a study of zeolites and activated carbons, it was proposed
that higher gas adsorption could be achieved on materials
with a large volume of micropores with an appropriate
diameter (although the specific ‘‘appropriate diameter’’ was
left undetermined).45 Much early work was directed toward
meeting the first part of this goal—namely, the synthesis of
highly porous frameworks, which could then be filled with
hydrogen gas. Some examples include several of the ‘‘MIL’’
series of MOFs reported by Férey et al., with pore sizes
greater than 25 Å,11 and the isostructural ‘‘IRMOF’’ series
with progressively larger pores reported by Yaghi et al.,7,12
all based on carboxylate ligands and possessing noninterpenetrated networks. These materials possess pore
volumes greater than 1.5 cm3 g21, and in turn adsorb a
sizable amount of hydrogen at high pressures: 6.01 wt% for
J. Mater. Chem., 2007, 17, 3154–3160 | 3155
MIL-10111 and 6.7 wt% for IRMOF-20 at 77 K.8 As much as
80% of the volume of these materials is empty space. In view of
Aristotle’s observation that ‘‘nature abhors a vacuum,’’ it is
unlikely that a material with a significantly greater permanent
porosity fraction can be created. Nonetheless, as shown in
Fig. 1, it is obvious that large pore volumes and surface
areas are necessary for high hydrogen saturation uptakes at
77 K; the same will be true for hydrogen storage at ambient
temperature.
Pore size and interpenetration
Various studies have reported the perhaps counterintuitive
finding that smaller pores actually take up hydrogen more
effectively than very large ones.16,37 The ideal pore size seems
to be 4.5–5 Å, or approximately 2.8–3.3 Å when the van der
Waals radii of the atoms composing the pore walls are
excluded; this is comparable to the y2.8 Å kinetic diameter of
H2. Pores of this size allow the dihydrogen molecule to interact
with multiple portions of the framework rather than just one
SBU or organic linker, increasing the interaction energy
between the framework and H2. Some have also proposed
that the curvature of pores or nonlinearity of channels, in
addition to the size or chemical functionality, plays a role,
although this is difficult to quantify experimentally.37,46
One problem encountered when attempting to quantify
the exact relationship between pore size and hydrogen
uptake is the variety of methods used to calculate and
Surface area, porosity, and hydrogen adsorption data for selected MOFs
Table 1
a
Material
Mn(HCO2)2
Mg3(HCO2)6
MOF-5, Zn4O(bdc)3, IRMOF-1
Sc2(bdc)3
MIL-53(Al), Al(OH)(bdc)
MIL-53(Cr), Cr(OH)(bdc)
MIL-101, Cr3OF(bdc)
IRMOF-2, Zn4O(bbdc)3
IRMOF-3, Zn4O(abdc)3
IRMOF-6, Zn4O(cbbdc)3
IRMOF-8, Zn4O(ndc)3
IRMOF-9, Zn4O(bpdc)3
IRMOF-11, Zn4O(hpdc)3
IRMOF-13, Zn4O(pydc)4
IRMOF-18, Zn4O(tmbdc)3
IRMOF-20, Zn4O(ttdc)3
Mg3(ndc)3
Mn(ndc)
Zn4O(L1)3
Zn4O(L2)3
Er2(pdc)3
Y2(pdc)3
MAMS-1, Ni8(tbbdc)6
MOF-74, Zn3O3(dhbdc)
HKUST-1, Cu3(btc)2
MIL-96, Al3O(btc)3
MIL-100, Cr3OF(btc)
Dy(btc)
TUDMOF-1, Mo3(btc)2
PCN-6, Cu3(tatb)2
PCN-9, Co4(tatb)8/3
MOF-177, Zn4O(btb)
MIL-102, Cr3OF(ntc)3/2
MOF-505, Cu2(bptc)
Cu2(tptc)
Cu2(qptc)
Zn3(bdt)3
Mn3(bdt)3
Mn3(bdt)8Cl2
Cu(bdt)
Mn(btt)
Pd(pymo)2
Cu(pymo)2
Zn7O2(pda)5
ZIF-8, Zn(mim)2
ZIF-11, Zn(pim)2
Zn(ndc)(bpe)K
Zn3(bpdc)3(bipy)
Co3(bpdc)3(bipy)
Ni(cyclam)(bpydc)
Zn2(bdc)(dabco)
b
2
SA /m g
21
Pore
volume/cm3 g21
d
240
150d
4170
721d
1590
1500
5500
2544
3062
3263
1818
2613
2340
2100
1501
4590
520
191
502d
396d
427d
676d
0.043
0.332
0.59
0.56
1.9
0.88
1.07
1.14
0.90
0.73
1.53
0.068
0.20
0.13
0.186
0.294
1132
2175
0.39
0.75
2800
655d
2010
3800
1355
5640
42.1
1830
2247d
2932d
640d
290d
530d
200f
2100d
600d
350d
1.0
1810
1947g
303
792e
922e
817
2090
0.67
1.453
0.51
0.12
0.680
0.886
1.138
0.795
0.17
0.663
0.582
0.2
0.33
0.38
0.37
0.75
3156 | J. Mater. Chem., 2007, 17, 3154–3160
H2 uptake at 77 K,
1 atm (wt%)
0.9
0.60
1.5
1.5
2.1
1.8
2.5
1.21
1.42
1.48
1.50
1.17
1.9
1.73
0.89
1.35
0.78
0.57
Maximum H2 uptake (wt%)
77 K
298 K
5.2, 48 bar
0.45, 60 bar
3.8, 16 bar
3.1, 15 bar
6.1, 40 bar
0.43, 80 bar
4.9, 32 bar
3.6, 10 bar
0.4, 30 bar
3.5, 34 bar
DHadsc/
kJ mol21
4.8
10i
6.1
9.0
6.7, 70 bar
9.5
1.12, 48 bar
0.98, 48 bar
0.675
0.760
0.6
1.77
2.54
1.6
1.0
1.32
1.75
1.9
1.53
1.25
0.65
2.59
2.52
2.24
1.4
0.9
0.8
0.66
2.1
1.2
0.8
2.3, 26 bar
3.6, 10 bar
1.96, 3 bar
3.28, 26.5 bar
0.35, 65 bar
8.3
6.8
0.15, 73.3 bar
6.3i
10.1
7.5, 70 bar
0.9, 10 bar
4.2h
6.7h
7.01h
0.05, 35 bar
5.99i
6.5j
8.7
8.4
8.8
6.9, 90 bar
1.4, 90 bar
10.1
1.01, 71.43 bar
1.29
1.37
0.8
1.74
1.98
1.1
2.01
3.1, 55 bar
2.0, 40 bar
0.3, 65 barh
7.1
6.8
Reference
4
5
6–8
9
10, 11
10, 11
11
7
7
7, 8
12, 13
7
7, 8
7
12
7, 8
6, 14
15
16
16
17
17
18
7, 8
7, 19
20
11
21
22
23, 24
25
8, 12
26
27–29
29
29
30
30
30
30
31
32
32
33
34
34
35
36, 37
36, 37
38
37
This journal is ß The Royal Society of Chemistry 2007
Fig. 1 Correlation between surface area (Langmuir method, N2) and
saturation hydrogen uptake at 77 K.
report the size of pores. One set of methods, the application of
Dubnin–Astakhov analysis or the Horvath–Kawazoe model to
gas sorption data, provides an estimation of pore size, but is
limited by the quality of the adsorption data and the gas used.
Given that MOFs typically have a highly ordered crystalline
structure, it is generally not difficult to determine a highquality single-crystal X-ray structure of these materials. Many
researchers apply a variety of software tools to estimate
accessible pore volume and pore sizes based on these structures; however, this method relies both on estimations of
the van der Waals radii of the atoms along the pore or
channel walls, and the particular algorithms of the software
packages used. We have found perhaps the most useful
information to be simple atom-to-atom distance measurements
across the pores or channels, as measured directly from the
crystal structure; this allows the reader or reviewer to make
‘‘apples-to-apples’’ comparisons from one paper to the next,
and to apply any additional assumptions or estimations as
necessary.
Interpenetration, the phenomenon in which two or more
networks in a structure are physically entangled (as links in a
chain), is a common motif in MOF structures.47 The typical
effect of interpenetration on porosity is to subdivide large
single pores, each bounded by the entire organic linker, into
several smaller ones, each bounded by smaller portions of the
organic linker. One extreme example of a highly interpenetrated structure with small pores is Zn4O(L1)3 (L1 =
6,69-dichloro-2,29-diethoxy-1,19-binaphthyl-4,49-dibenzoate),16
with a fourfold interpenetrating structure, open channels of
less than 5 Å, and BET surface area of only 502 m2 g21; this
material adsorbs 1.12 wt% of hydrogen at room temperature
and 48 bar, as contrasted to MIL-101, with a surface area of
5500 m2 g21 and hydrogen uptake of only 0.43 wt% under the
same conditions.
We have recently reported a system which can be
synthesized in either an interpenetrated or a noninterpenetrated form, allowing evaluation of interpenetration as an
independent criterion resolved from other factors.23 This
MOF, Cu3(tatb)2 (also designated PCN-6), contains the
copper-carboxylate paddlewheel SBU linked by a trigonal
triazine-based ligand, and the catenation isomerism is controlled by the presence or absence of oxalic acid (apparently in
a templating role, as it does not appear in the final structure).
Not surprisingly, the noninterpenetrated form (PCN-69) has a
Table 1 Surface area, porosity, and hydrogen adsorption data for selected MOFs (Continued )
a
Material
b
2
21
SA /m g
Pore
volume/cm3 g21
H2 uptake at 77 K,
1 atm (wt%)
Maximum H2 uptake (wt%)
77 K
298 K
DHadsc/
kJ mol21
Reference
Ni2(dhtp)
1083
0.41
0.7
1.8, 60 bar
0.3, 65 bar
39
700d
0.94
10.4
40
NaNi3(OH)(sip)
0.181
0.8
41
Ni2(bipy)3(NO3)4
0.63
0.7
2.5, 15 bar
0.15, 15 bar
41
Ni3(btc)2(pic)6(pd)
1670
0.59
2.08
37
Zn2(bdc)(tmbdc)(dabco)
1400
0.50
1.85
37
Zn2(tmbdc)2(dabco)
1450
0.52
1.70
37
Zn2(ndc)2(dabco)
1610
0.57
1.78
37
Zn2(tfdbc)2(dabco)
1740
0.62
1.68
37
Zn2(tmbdc)2(bipy)
IRMOF-8 + Pt/AC
3.5, 100 bar
24.8
42
a
Abbreviations: bdc = 1,4-benzenedicarboxylate, bbdc = 2-bromo-1,4-benzenedicarboxylate, abdc = 2-amino-1,4-benzenedicarboxylate, cbbdc =
1,2-cyclobutane-3,6-benzenedicarboxylate, ndc = 2,6-naphthalenedicarboxylate, bpdc = 4,49-biphenyldicarboxylate, hpdc = 4,5,9,10tetrahydropyrene-2,7-dicarboxylate, pydc = pyrene-2,7-dicarboxylate, tmbdc = 2,3,5,6-tetramethylbenzene-1,4,-dicarboxylate, ttdc = thieno[3,2b]thiophene-2,5-dicarboxylate, L1 = 6,69-dichloro-2,29-diethoxy-1,19-binaphthyl-4,49-dibenzoate, L2 = 6,69-dichloro-2,29-benzyloxy1,19-binaphthyl-4,49-dibenzoate, pdc = pyridine-3,5-dicarboxylate, tbbdc = 5-tert-butyl-1,3-benzenedicarboxylate, dhbdc = 2,5-dihydroxy-1,4benzenedicarboxylate, btc = 1,3,5-benzenetricarboxylate, tatb = triazine-4,49,40-s-triazine-2,4,6-triyltribenzoate, btb = 1,3,5-benzenetribenzoate,
ntc = naphthalene-1,4,5,8-tetracarboxylate, bptc = biphenyl-3,39,5,59-tetracarboxylate, tptc = terphenyl-3,30,5,50-tetracarboxylate, qptc =
quaterphenyl-3,3-9,5,5--tetracarboxylate, bdt = 1,4-benzeneditetrazolate, btt = 1,3,5-benzenetristetrazolate, pymo = 2-pyrimidinolate, pda = 1,4phenylenediacrylate, mim = methylimidizolate, pim = phenylimidizolate, bpe = trans-1,2-bis(4-pyridyl)ethene, bipy = 2,29-bipyridine, cyclam =
1,4,8,11-tetraazacyclotetradecane, bpydc = 2,29-bipyridyl-5,59-dicarboxylate, dabco = 1,4-diazabicyclo[2.2.2]octane, dhtp = 2,5dihydroxyterephthalate, sip = 5-sulfoisophthalate, pic = 3-picoline, pd = 1,2-propanediol, tmbdc = 2,3,5,6-tetramethylbenzene-1,4,-dicarboxylate,
tfbdc = 2,3,5,6-tetrafluoro-1,4-benzenedicarboxylate. b Surface area determined using Langmuir method, N2 adsorption, 77 K, unless otherwise
noted. c At low coverage, calculated from adsorption isotherms at two or more temperatures, unless otherwise noted. d BET method, N2
adsorption, 77 K. e BET method, Ar adsorption, 87 K. f BET method, O2 adsorption, 77 K. g Calculated from crystal structure. h Estimated
saturation limit from Langmuir plot. i Determined directly by microcalorimetry. j Calculated from Grand Canonical Monte Carlo simulation.
This journal is ß The Royal Society of Chemistry 2007
J. Mater. Chem., 2007, 17, 3154–3160 | 3157
higher overall porosity, based on the solvent-accessible volume
calculated from the single-crystal X-ray structure. However,
the interpenetrated form of this MOF exhibits a 41% increase
in surface area, 133% increase in volumetric hydrogen uptake,
and 29% increase in gravimetric hydrogen uptake when
compared to the noninterpenetrated form.23
Two Grand Canonical Monte Carlo (GCMC) studies on
interpenetrated MOFs in the IRMOF series demonstrated that
in these structures, the smaller pore size and multiplicity of
networks allows the dihydrogen molecule to interact with the
central (phenyl-containing) portion of multiple ligands, thus
increasing the relevance of the non-coordinating portion of the
ligand to hydrogen storage.28,48 Both studies agree that the
importance of a high heat of adsorption is greatest at low
loadings (low H2 pressures), and that the overall pore volume
becomes more important at higher loadings. The net result
of interpenetration is an increase in the interaction energy,
DHads, and is reflected in an increase in the H2 uptake at 77 K
and 1 atm.
Thermal activation and metal-based H2 binding sites
In addition to solvent molecules trapped within the pores of
the material, as-synthesized MOFs may also have solvent
molecules attached as ligands to the metal centers or
incorporated as part of the SBU. Removal of these solvent
ligands is referred to as thermal activation, and is often a
necessary step to access the full gas-adsorption potential of a
material. In some cases, these coordinated solvent molecules
may merely protrude into windows or channels, blocking
access of dihydrogen molecules into the larger spaces within
the framework. In other cases, removal of these ligands
(often aqua ligands) leaves the metal cation exposed on
the interior surfaces and open to direct approach by the
dihydrogen molecule.
Chen and coworkers have shown that the removal of axial
aqua ligands from dicopper paddlewheel SBUs via thermal
activation exposes the copper binding sites in MOF-505;27
Long et al. have demonstrated the same phenomenon in a
Mn-containing MOF with tetrazolate ligands,30 as have
Bordiga et al. in HKUST-1.49 This generates a so-called
unsaturated metal center (UMC), which then interacts strongly
with the dihydrogen molecule. These UMCs can be seen as
analogous to entatic metal centers in bioinorganic chemistry,
in which metal ions (such as the iron in hemoglobin) are
forced into an unusual coordination geometry (see Fig. 2)—
such a concept has been advanced in our lab by the study
of PCN-9, a MOF containing a coordinatively unsaturated
Co4(m4-O)(CO2)8 SBU.25 This study also demonstrated
via IR spectroscopy that the cobalt site can bind probe
molecules CN2 and CO;25 similar results were found in CO
binding to the copper center of HKUST-1.49 Exposing these
sites greatly increases the ability of the material to adsorb
hydrogen; the incorporation of accessible UMCs has been
shown to be a viable strategy to increase the hydrogen uptake
by MOFs.
Single-crystal neutron diffraction of MOF-5 revealed two
hydrogen-binding sites, one higher-energy site over the center
of the Zn4(m4-O)(CO2)8 SBU, and a second site over the face of
3158 | J. Mater. Chem., 2007, 17, 3154–3160
Fig. 2 (a) Schematic drawing of the active center of hemoglobin; the
gold sphere represents an iron atom. (b) Schematic drawing of
the tetranuclear M4(m4-O)(CO2)8 SBU containing coordinatively
unsaturated metal atoms. (c) The M2(CO2)4 paddlewheel SBU.
a ZnO4 tetrahedron.50 Neutron powder diffraction reveals two
additional sites in MOF-5 at increased loading: one associated
with the zinc-carboxylate moiety Zn(CO2), and one over the
phenyl ring of the ligand.51 Generally these agree with an
inelastic neutron scattering experiment performed on the same
material, differing only in the preferred order of site
occupation. MP2 ab initio calculations of simplified MOF
models predict the metal-carboxylate linker to be a highenergy binding site, with additional lower-energy sites located
around the phenyl ring of the ligand.52
Additional studies probing H2 binding sites have been
performed on a number of other MOFs as well. Long et al.
have employed neutron powder diffraction to confirm that H2
is closely associated with the UMCs found in a Mn-tetrazole
MOF.31 A combination of temperature-programmed desorption and inelastic neutron scattering of H2-loaded NaNi(sip)2
(sip = 5-sulfoisophthalate) has revealed a number of discrete
H2 binding sites, the strongest of which can be associated with
an unsaturated Ni site.40 A combined DFT and GCMC
dynamical study of the dicopper-paddlewheel-containing
MOF-505 shows that, as expected based on the earlier
thermal-activation study (vide supra), the binding energy of
H2 on MOF-505 is highest at the copper UMC sites exposed
by thermal activation.27
Most recently, neutron powder diffraction of D2-loaded
HKUST-1 identified six distinct D2 sites, shown in Fig. 3. The
first, highest-energy site is associated with the copper UMCs
(on the axes of the paddlewheel SBUs), near enough to
indicate significant interaction with the d9 Cu(II) center. The
remaining sites fill competitively from the smallest to largest
pores, with these sites located near the benzene ring and
carboxylate moieties of the ligand.53
Ligand structure and functionalization
Although the metal sites and/or the SBU are the preferential
adsorption sites for hydrogen, the organic linker can play an
important secondary role in increasing adsorption further, as
shown above. In the IRMOF series developed by Yaghi et al.,
the basic structural motif of Zn4(m4-O)(CO2)8 SBUs connected
by aromatic phenyl-containing linkers is repeated to generate a
series of isostructural materials, which differ only in the central
portion of the ligand.7,12 Increasing the aromaticity of this
central portion, from a simple phenyl ring (MOF-5/IRMOF-1)
to cyclobutylbenzene (IRMOF-6) to naphthalene (IRMOF-8)
This journal is ß The Royal Society of Chemistry 2007
Dissociative adsorption of hydrogen
Fig. 3 D2 Sites in HKUST-1, identified via neutron powder diffraction, numbered in order of occupation with increased loading. Top:
shown along [001] (left) and [111] (right). Bottom: axial Cu(II)
paddlewheel UMC site (left), along [111] in the 5 Å small pore with
3.5 Å side windows (middle), and along [100] showing the 9 Å pore.
Reprinted with permission from ref. 53. Copyright 2006, American
Chemical Society.
One potential method for enhancing the hydrogen uptake
involves the concept of ‘‘hydrogen spillover’’. Long known on
supported metal catalysts, spillover occurs by the dissociative
chemisorption of hydrogen onto the metal surface, followed
by the migration of monoatomic hydrogen onto the supporting material.56 Intimate contact between this supporting
material and a third material allows ‘‘secondary spillover’’;
this technique has been applied by Li et al. in preliminary
studies, enhancing the hydrogen adsorption of MOF-5 and
IRMOF-8 as much as eightfold.42 To achieve this remarkable
result, 5% Pt on activated carbon was mixed with the MOF
and sucrose, then ground to make an intimate mixture; melting
and subsequent carbonization of the sucrose provided a
carbon ‘‘bridge’’ permeating the mixture, allowing atomic
hydrogen to migrate from the metal surface across the carbon
to the MOF. The smaller size of the hydrogen atom should
allow the adsorbate to penetrate smaller pores which are
inaccessible to the dihydrogen molecule; in other words, the
potential exists to bypass the gravimetric storage limits
imposed by the density and incompressibility of liquid
dihydrogen. Dissociative adsorption should also increase the
affinity of the framework for hydrogen, as the adsorbed
species in this case is the hydrogen atom rather than the
dihydrogen molecule.
Conclusions
increases the hydrogen uptake dramatically, from 0.5 wt% to
1.0 wt% and 1.5 wt%, respectively.12
One further way to enhance the affinity of the ligand for
the dihydrogen molecule is by chemical functionalization: the
introduction of an electron-donating group (or groups) to
the central portion of the ligand. Again, the IRMOF series
provides illustration: replacing one hydrogen atom on the
central benzene ring of the linker in MOF-5 with –Br or –NH2
affords IRMOF-2 and -3, respectively; replacing all four
hydrogens with methyl groups affords IRMOF-18.7 A fifth
member of the series, IRMOF-20, replaces the phenyl ring
of bdc with a thieno-[3,2-b]thiophene moiety; this was
expected to enhance hydrogen adsorption significantly due to
the increased polarizability of the heteropolycyclic ligand.7
MP2 computational studies suggest that these electrondonating substituents should increase the affinity of the
aromatic ring for H2;54 however, little enhancement was found
in these functionalized IRMOFs in practice.7 A similar lack of
enhancement was reported by Kim et al. for pillared MOFs
constructed in part by ligands with all phenyl H atoms
replaced with either –F or –CH3.37 This may be due to partial
restriction of the pore size or blocking of the high-affinity
metal-based binding sites by the larger ligand, effectively
canceling out any benefit derived from electronic enhancement
of the ligand. It has also been proposed that cyanide and
N-heterocyclic ligands in general may have a higher affinity for
dihydrogen than purely graphitic ligands, in part based on
hydrogen adsorption studies with carbon, carbon nitride, and
boron nitride nanotube structures;55 MOFs that exhibit this
functionality include systems using triazine- and tetrazolatebased ligands.24,30
This journal is ß The Royal Society of Chemistry 2007
In order to reach hydrogen storage goals, a method must be
found which has both sufficient capacity and acceptable
uptake/release kinetics and thermodynamics at ambient
temperatures and reasonable operating pressures. The density
of molecular dihydrogen imposes strict limits on the potential
of tank storage technologies to meet the 2015 gravimetric
goals; hydride species uptake sufficient amounts of atomic
hydrogen, but the thermodynamics of these systems raises
daunting challenges. Current MOFs and other porous
physisorbents present moderate uptake volumes but insufficient interaction energy to retain hydrogen at ambient
temperature. However, MOFs offer a path to the middle
ground—the volume advantages of tanks and porous
adsorbents and the potential of moderate affinity to hydrogen,
allowing uptake and release at reasonable temperatures.
Although the heat of adsorption of dihydrogen on current
MOFs is still quite low, this property appears ‘‘tunable’’ in
MOFs by a variety of methods, potentially allowing these
materials to reach a ‘‘quasi-chemisorptive’’ regime in which a
higher-energy interaction between dihydrogen and the framework exists but remains short of true chemical bonding.
Catalyzed dissociation allows the material to interact with
atomic hydrogen rather than molecular dihydrogen, increasing
further both the interaction energy and the volume fraction
which can be penetrated by the adsorbed hydrogen. The
next several years are sure to see both an increased
understanding of the tunability of these interactions and
the further development of novel MOFs with ever-greater
uptake capacities—all in the hopes of a more efficient, cleaner
energy future.
J. Mater. Chem., 2007, 17, 3154–3160 | 3159
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