A robust, air-stable, reusable ruthenium catalyst for dehydrogenation of ammonia borane

Article abstract of DOI:10.1021/ja2058154

We describe an efficient homogeneous ruthenium catalyst for the dehydrogenation of ammonia borane (AB). This catalyst liberates more than 2 equiv of H₂ and up to 4.6 system wt % H₂ from concentrated AB suspensions under air. Importantly, this catalyst is robust, delivering several cycles of dehydrogenation at high [AB] without loss of catalytic activity, even with exposure to air and water.

Full text of DOI:10.1021/ja2058154

COMMUNICATION
pubs.acs.org/JACS
A Robust, Air-Stable, Reusable Ruthenium Catalyst for
Dehydrogenation of Ammonia Borane
Brian L. Conley, Denver Guess, and Travis J. Williams*
Department of Chemistry and Loker Hydrocarbon Institute, University of Southern California, 837 Bloom Walk, Los Angeles,
California 90089-1661, United States
S Supporting Information
b
Scheme 1. Dehydrogenation of Ammonia Borane to Produce
Hydrogen, Borazine, and Polyborazylene
ABSTRACT: We describe an efficient homogeneous ruthe-
nium catalyst for the dehydrogenation of ammonia borane
(AB). This catalyst liberates more than 2 equiv of H2 and up
to 4.6 system wt % H2 from concentrated AB suspensions
under air. Importantly, this catalyst is robust, delivering
several cycles of dehydrogenation at high [AB] without loss
of catalytic activity, even with exposure to air and water.
atalytic dehydrogenation of ammonia borane (AB) is an
C
active research area because AB is a potentially useful
hydrogen storage medium for transportation.1 This is because
of its high hydrogen density (19.6 wt %), ability to release H2
under mild conditions (both thermal2 and catalytic), and its desir-
able physical properties. Though catalytic hydrolysis is well-
known and very efficient for H2 production from AB,1b anhy-
drous dehydrogenation can enable a more efficient fuel cycle.
This is because hydrolysis reactions form stoichiometric quan-
tities of ammonia, a hydrogen fuel cell poison, and strong BꢀO
bonds, which preclude an efficient regeneration scheme.3 Several
heterogeneous4 and homogeneous5 transition-metal catalysts are
active for AB dehydrogenation, but these are limited by protic
and oxidative decomposition in air, low extent of H2 release
(e2 equiv), uncontrolled rate of H2 release, or production of
unwanted products such as ammonia (NH3) or borazine
(N3B3H6), which are poisonous to fuel cells. Additionally,
many are not viable because the catalyst loading is high or the
catalyst is not reusable. We report here a system that is long-
lived [turnover number (TON) > 5000], functions under air,
and dehydrogenates AB to give >2 equiv of H2. These charac-
teristics make it a leading candidate for use in a commercial
H2 storage system.
Figure 1. (left) H2 production from AB with 5.0 mol % 1 under air at
70 °C in 11:1 diglyme/benzene solution. (right) H2 production from AB
in the presence of 2.0 mol % 1 in tetraglyme under air at 70 °C at high
[AB]. H2 release = 4.2 system wt %.
dehydrogenation8 liberates only 1.6 equiv of H2 at 70 °C at a
rate that is ca. 2.5-fold slower under our reaction conditions (see
the Supporting Information).
This catalyst system is also remarkably reusable. Data for
successive runs at 2.0 mol % loading (70 °C) indicate that the
catalysis is efficient through four ca. 6 h runs at high [AB],
although the reaction slurry becomes viscous. To help alleviate
this situation, we added 0.2 mL of tetraglyme while recharging
the reactor with AB for subsequent runs. In this way, we ran four
successive runs in a single reactor. The rates for all of these were
similar, and they released 2.2, 2.1, 2.3, and 2.2 equiv of H2,
respectively, for a total TON of 440. With a catalyst loading of 0.1
mol %, we observed a TON of ca. 5700 over three runs and
liberation of up to 4.6 wt % H2.9 For comparison, Fagnou
reported a ruthenium catalyst that could release up to 1.0 system
wt % H2 from AB at room temperature and up to 3.6 system wt %
Catalyst 16 liberates H2 from AB under air in diglyme solu-
tions (Scheme 1). To determine the efficiency of AB dehydro-
genation with 1, we treated a 0.42 M AB solution with 5.0 mol %
1 at 70 °C in a Schlenk flask under air and monitored the
production of H2 using a eudiometer. The system produced 1.94
equiv of H2 in 7 h (Figure 1 left). Added metallic mercury did not
inhibit the reaction under these conditions, which indicates
homogeneous catalysis.7 We also examined more practical,
highly concentrated suspension conditions. A slurry of AB and
tetraglyme (100.0 mg of AB, 202.6 mg of tetraglyme, 5.8 system
wt % stored H2) was treated with 2.0 mol % 1 at 70 °C and
yielded 2.0 equiv of H2 in 4 h and a total of 2.2 equiv (TON =
110) upon completion of the reaction. This corresponds to
the release of 4.2 wt % H2. In contrast, uncatalyzed AB
Received: June 26, 2011
Published: August 09, 2011
r
2011 American Chemical Society
14212
dx.doi.org/10.1021/ja2058154 J. Am. Chem. Soc. 2011, 133, 14212–14215
|

Journal of the American Chemical Society
COMMUNICATION
H2 from AB/MeNH2BH3 mixtures at 50 °C.5b This is the best
previously published weight-content H2 release for a homoge-
neous catalytic system dehydrogenating an amineborane.
Our system works well when open to air, which is beneficial for
commercialization: a system stable to air and water provides
practical advantages over a more sensitive one. Measurement of
H2 produced in eudiometry studies in air and under N2 were
identical under concentrated AB conditions. To quantify this
effect further, we compared rate constants (kobs) for [AB]
consumption obtained for reactions conducted under nitrogen
and air. In this experiment, 2.5 mol % 1 was prepared in diglyme/
benzene-d6 in a glovebox and sonicated open to air for 1 h prior
to addition of AB to ensure removal of the N2 environment and
establish atmospheric O2 levels in the NMR tube. The kobs values
at 70 °C were 2.65(9) ꢁ 10ꢀ4 sꢀ1 in air and 2.54(9) ꢁ 10ꢀ4 sꢀ1
in N2. Thus, we observed neither inhibition nor acceleration of
the catalysis. This suggests that at ambient levels of O2, the
catalysis is not shut down by decomposition of the active species
or accelerated by formation of a more reactive, higher-valent
ruthenium oxo complex.
We probed the mechanism of AB dehydrogenation with 1 by
11B NMR spectroscopy in dilute diglyme solutions. Unlike
Shvo’s catalyst,5e 1 demonstrates first-order kinetics in [AB]
through 3 half-lives, as observed by the disappearance of the AB
resonance in the 11B NMR spectrum. Similar to others’ observa-
tions, the intermediates we observed in these reactions are
branched and unbranched cyclotriborazane, aminoborane oligo-
mers, borazine, and polyborazylene (PB),1a,1d,8 with the latter
two being the major products upon completion. As is the case for
Shvo’s catalyst,5e we did not observe the formation of insoluble B,
N oligomers in this system.
The appearance of PB at early times in reactions with 1 is
encouraging, whereas none is seen under Shvo conditions.
Formation of PB is a feature of extensive dehydrogenation,
indicating the production of more than 2 equiv of H2, and it
has been proposed to occur by borazine cross-linking in transi-
tion-metal-catalyzed systems. However, in contrast to (NHC)Ni
systems, in which borazine is quickly consumed,5d [borazine]
increases over time in reactions of 1 and is not consumed late in
the reaction.10 This suggests that catalytic borazine cross-linking
is not the mechanism for PB formation in this system. We suspect
that the operative mechanism involves reaction of borazine with
AB or one of its early dehydrogenation products. Efficient,
catalytic borazine cross-linking has not been demonstrated to
date in the absence of other B,N materials.11
Scheme 2. Synthesis of Bridging Carboxylate Complexes 2
and 3a
a (left) Synthesis. (right) ORTEP plot of 3 with 50% probability
ellipsoids. H atoms and the triflate counterion have been omitted for
clarity.
of AB.14 The kH/kD values for the BꢀH and NꢀH bonds were
1.22(14) and 1.58(9), respectively. A combined isotope effect
measured using D3NꢀBD3 (kNHBH/kNDBD) was 1.67(18).
Thus, the product of the two independently measured isotope
effects [1.92(25)] is within experimental error of the measured
double isotope effect, which is consistent with a concerted,
asynchronous transition state in the rate-determining step. How-
ever, this conclusion is tentative because we cannot be certain
whether the observed isotope effects involve H2 transfer from AB
to the catalyst or multiple steps, as in the (NHC)Ni system.13a
Other mechanistic scenarios have the bridging hydroxyl group
of 1 donating H+ to the solution, thus initiating an acid-catalyzed
reaction,15 or the thus-formed oxide bridge providing an internal
base in a bifunctional mechanism. To probe these, we synthe-
sized bridging carboxylate complexes 2 and 3 (Scheme 2 left),
which are devoid of the protic functionality. The crystal structure
of 3 (Scheme 2 right)16 shows an unprecedented RuꢀB μ2-kO:k
O0 bonding mode for the trifluoroacetate wherein the boron is
distorted from its tetrahedral geometry to give a trigonal-
pyramidal structure. This behavior could be rationalized by
unfavorable steric interactions between the bridging carboxylate
and other ligands in the unobserved μ2-kO:kO (single bridging
oxygen) mode. We suspect the same for 2.
Surprisingly, 5.0 mol % solutions of 1 and 2 have the same rate
constant for AB consumption [kobs = 3.74(7) ꢁ 10ꢀ4 and
Rate data (kobs, 70 °C) collected by 11B NMR spectroscopy
were pseudo-first-order in AB. A catalyst order study was
conducted by comparing the kobs values measured at various
concentrations of 1. The data gave a linear plot of ln(kobs) versus
ln([1]) with a slope of 0.56, which indicates the presence of a
dinuclear (Ru)2 intermediate,12 although we did not observe this
species. Possible structures include those containing a BꢀOꢀB
bridge, a bridging hydride, or a product of AB dehydrogenation
serving as the bridging moiety. Along these lines, we did not
observe a persistent metal hydride, which argues against a RuꢀH
moiety as the resting state of the catalyst. This is in line with
observations of the (NHC)Ni systems5d but contrasts our
own5e and others’ findings for ruthenium, and iridium5a catalysts.
We suspect that 1 might be analogous to several known
catalysts that interact with the polarized BꢀH and NꢀH bonds
of AB concurrently in a bifunctional transition state.5b,5f,13 We
probed this by recording deuterium isotope effects for isotopologues
3.72(10) ꢁ 10ꢀ4
s
ꢀ1, respectively, at 70 °C]. This disfavors a
H+-catalyzed mechanism. To interrogate the potential role of
boron as a Lewis acid, we attempted AB dehydrogenation with
complex 3, which has different electronic characteristics than 1 or
2. This gives faster catalysis: 2.5 mol % solutions of 1 and 3 gave
rate constants of kobs = 2.66(12) ꢁ 10ꢀ4 and 5.77(12) ꢁ 10ꢀ4 sꢀ1
,
respectively, for AB consumption at 70 °C. Thus, the occupancy of
the bridging coordination site between Ru and B has an important
influence on the dehydrogenation rate.
In multiple solvents we observed that the acetate and trifluor-
oacetate complexes hydrolyze completely in the presence of trace
water to re-form bridging hydroxide complex 1 and the corre-
sponding carboxylic acid, which indicates a thermodynamic
preference for the Ru(OH)B bridge. Adding water (ca. 6 equiv)
to a sample of 3 in tetrahydrofuran-d8 reveals that the hydrolysis
has a half-life of 6ꢀ7 h at room temperature, thus demonstrating
the lability of the carboxylate ligand. Furthermore, a methylene
14213
dx.doi.org/10.1021/ja2058154 |J. Am. Chem. Soc. 2011, 133, 14212–14215

Journal of the American Chemical Society
COMMUNICATION
chloride-d2 solution of 3 and 1 equiv of TFA-d1 showed the
coalescence of the bound and free trifluoroacetic acid (TFA)
signals at 50 °C as monitored by 19F NMR spectroscopy: two
fluorine resonances [δ ꢀ75.38 (bound) and ꢀ75.25 (free) at
25 °C] coalesced to one (δ ꢀ75.51). This observation indicates
that exchange occurs rapidly on the NMR time scale. Thus, the
bridging trifluoroacetate appears to be sufficiently labile to allow
boron and ruthenium to participate in catalysis, but these data do
not rule out participation of the acetate and trifluoroacetate in the
mechanism of AB dehydrogenation by 2 and 3, as has been
documented in systems for hydrocarbon CꢀH activation.17
Importantly, dehydrogenation is not efficient in the absence of
the borate ligand. Catalyst 1’s nonligated synthetic precursor,
[(cym)RuCl2]2 (4), does not participate in efficient catalysis;
these reactions precipitate metallic material under our experi-
mental conditions. 11B NMR data for catalysis with 4 revealed a
rate of AB consumption that is only ca. 2-fold lower than that of
1; however, H2 production from this system was limited to
1 equiv, and it is not reusable. Additionally, added metallic
mercury significantly attenuated the production of H2 with 4.
Full kinetics data for AB dehydrogenation with 4, Na[(2-py)2-
BMe2], and [((2-py)2BMe2)RuCl(cym)] (5) are described in
the Supporting Information.
In conclusion, we have demonstrated an efficient and robust
catalyst for highly productive ammonia borane dehydrogenation.
This catalyst liberates up to 4.6 wt % H2 from AB suspensions
and is resistant to deactivation in air, which makes it one of the
most appealing homogeneous transition-metal catalysts de-
signed to date. Its longevity at low catalyst loadings (TON up
to 5700) and its air stability are unprecedented in transition-
metal-catalyzed AB dehydrogenation. Initial mechanistic inves-
tigations based on (1) isotope effects and (2) the relative reaction
rates of 2 and 3 suggest that dual-site cooperativity could be
operative. This mechanistic insight may lead to the development
of more efficient systems for AB dehydrogenation. Ongoing
work in our laboratory involves uncovering the detailed roles of
the boron and ruthenium centers in AB dehydrogenation and the
application of dual-site catalysis to more general hydride manip-
ulation reactions.
’ REFERENCES
(1) (a) Staubitz, A.; Robertson, A. P. M.; Manners, I. Chem. Rev.
2010, 110, 4079–4124. (b) Stephens, F. H.; Pons, V.; Baker, R. T. Dalton
Trans. 2007, 2613–2626. (c) Marder, T. B. Angew. Chem., Int. Ed. 2007,
46, 8116–8118. (d) Hamilton, C. W.; Baker, R. T.; Staubitz, A.;
Manners, I. Chem. Soc. Rev. 2009, 38, 279–293 and references therein.
(2) (a) Baitalow, F.; Bauman, J.; Wolf, G.; Jaenicke-Rossler, K.;
Leitner, G. Thermochim. Acta 2002, 391, 159–168. (b) Wolf, G.;
Baumann, J.; Baitalow, F.; Hoffmann, F. P. Thermochim. Acta 2000,
343, 19–25. (c) Wang, J. S.; Geanangel, R. A. Inorg. Chim. Acta 1988,
148, 185–190. (d) Bluhm, M. E.; Bradley, M. G.; Butterick, R., III;
Kusari, U.; Sneddon, L. G. J. Am. Chem. Soc. 2006, 128, 7748–7749. (e)
Rassat, S. D.; Aardahl, C. L.; Autrey, T.; Smith, R. S. Energy Fuels 2010,
24, 2596–2606.
(3) (a) Davis, B. L.; Dixon, D. A.; Garner, E. B.; Gordon, J. C.; Matus,
M. H.; Scott, B. Angew. Chem., Int. Ed. 2009, 48, 6812–6816. (b) Sutton,
A. D.; Burrell, A. K.; Dixon, D. A.; Garner, E. B., III; Gordon, J. C.;
Nakagawa, T.; Ott, K. C.; Robinson, J. P.; Vasiliu, M. Science 2011,
331, 1426–1429.
(4) (a) Jaska, C. A.; Temple, K.; Lough, A. J.; Manners, I. J. Am.
Chem. Soc. 2003, 125, 9424–9434. (b) Jaska, C. A.; Manners, I. J. Am.
Chem. Soc. 2004, 126, 1334–1335. (c) Shrestha, R. P.; Diyabalanage,
H. V. K.; Semelsberger, T. A.; Ott, K. C.; Burrell, A. K. Int. J. Hydrogen
Energy 2009, 34, 2616–2621.
(5) (a) Denney, M. C.; Pons, V.; Hebden, T. J.; Heinekey, M.;
Goldberg, K. I. J. Am. Chem. Soc. 2006, 128, 12048–12049. (b) Blaquiere,
N.; Diallo-Garcia, S.; Gorelsky, I.; Black, A.; Fagnou, K. J. Am. Chem. Soc.
2008, 130, 14034–14035. (c) K€ass, M.; Friedrich, A.; Drees, M.;
Schneider, S. Angew. Chem., Int. Ed. 2009, 48, 905–907. (d) Keaton,
R. J.; Blacquiere, M.; Baker, R. T. J. Am. Chem. Soc. 2007, 129,
1844–1845. (e) Conley, B. L.; Williams, T. J. Chem. Commun. 2010,
46, 4815–4817. (f) Chapman, A. M.; Haddow, M. F.; Wass, D. F. J. Am.
Chem. Soc. 2011, 133, 8826–8829. (g) Wright, W. R. H.; Berkeley, E. R.;
Alden, L. R.; Baker, R. T.; Sneddon, L. G. Chem. Commun. 2011,
47, 3177–3179. (h) Kim, S.-K.; Han, W.-S.; Kim, T.-J.; Kim, T.-Y.; Nam,
S. W.; Mitoraj, M.; Piekoꢀs, y.; Michalak, A.; Hwang, S.-J.; Kang, S. O.
J. Am. Chem. Soc. 2010, 132, 9954–9955 .
(6) Conley, B. L.; Williams, T. J. J. Am. Chem. Soc. 2010, 132,
1764–1765.
(7) (a) Anton, D. R.; Crabtree, R. H. Organometallics 1983,
2, 855–859. (b) Widegren, J. A.; Finke, R. G. J. Mol. Catal. A: Chem.
2003, 198, 317–341. The mercury drop test is commonly invoked to
provide positive evidence for homogeneous catalysis when a reaction is
not inhibited upon its addition. As pointed out by a reviewer and
thoroughly discussed in part (b) of this reference, a Hg drop test is not
definitive. Some transition metals (e.g., Ru) are more difficult to
amalgamate, and care should be taken in analysis when using precursors
based on these metals. More informative to the point of homogeneous vs
heterogeneous catalysis, though not conclusive, are kinetic behavior,
solution turbidity, and order with respect to precatalyst. In our case, the
sum of our results suggests homogeneous behavior.
’ ASSOCIATED CONTENT
S
Supporting Information. Full experimental procedures,
b
characterization of complexes 2 and 3, full kinetics data for AB
dehydrogenation with 1ꢀ5 and Na[(2-py)2BMe2], a comment
on the [EtOH] dependence of the rate, and crystallographic data
(CIF) for 3. This material is available free of charge via the
Internet at http://pubs.acs.org.
(8) Shaw, W. J.; Linehan, J. C.; Szymczak, N. K.; Heldebrant, D. J.;
Yonker, C.; Camaioni, D. M.; Baker, R. T.; Autrey, T. Angew. Chem., Int.
Ed. 2008, 47, 7493–7496.
’ AUTHOR INFORMATION
(9) System wt % H2 production was calculated over four cycles
(combined) using the following formula: [(theoretical wt yield of H2)/
(total wt of AB + wt of solvent + wt of catalyst)] ꢁ (total equiv of
liberated H2/theoretical equiv of H2) ꢁ 100%. For the four cycles at 2.0
mol %, this gave [(78.4 mg)/(400.0 mg + 810.4 mg + 38.0 mg)] ꢁ (8.8
equiv/12.0 equiv) ꢁ 100% = 4.6 system wt %.
Corresponding Author
travisw@usc.edu
’ ACKNOWLEDGMENT
(10) Slow thermal cross-linking was observed, as described by:
Fazen, P. J.; Remsen, E. E.; Beck, J. S.; Carroll, P. J.; McGhie, A. R.;
Sneddon, L. G. Chem. Mater. 1995, 7, 1942–1956.
We thank the National Science Foundation (CHE-1054910),
the University of Southern California, the Loker Hydrocarbon
Research Institute, and the Hydrocarbon Research Foundation
for research support and the NSF (DBI-0821671, CHE-
0840366) and NIH (S10-RR25432) for NMR spectrometers.
We thank Ralf Haiges for the X-ray study of 3.
(11) Baker and co-workers5d reported extensive dehydrogenation of
AB to polyborazylene, but control reactions with added borazine and
preformed catalyst revealed no cross-linking. We also failed to observe cross-
linking of borazine with 1 under the reaction conditions reported herein.
14214
dx.doi.org/10.1021/ja2058154 |J. Am. Chem. Soc. 2011, 133, 14212–14215

Journal of the American Chemical Society
COMMUNICATION
(12) Espenson, J. H. Chemical Kinetics and Reaction Mechanisms, 2nd
ed.; McGraw-Hill: New York, 2002; pp 77ꢀ82.
(13) (a) Zimmerman, P. M.; Paul, A.; Zhang, Z.; Musgrave, C. B.
Angew. Chem., Int. Ed. 2009, 48, 2201–2205. (b) Yang, X.; Hall, M. B.
J. Am. Chem. Soc. 2008, 130, 1798–1799.
(14) Hu, M. G.; Van Paasschen, J. M.; Geanangel, R. A. J. Inorg. Nucl.
Chem. 1997, 39, 2147–2150.
(15) Stephens, F. H.; Baker, R. T.; Matus, M. H.; Grant, D. J.; Dixon,
D. A. Angew. Chem., Int. Ed. 2007, 46, 746–749.
(16) CCDC 830936 contains the supplementary crystallographic
data for compound 3. These data can be obtained free of charge from
The Cambridge Crystallographic Data Centre via www.ccdc.cam.ac.uk/
data_request/cif.
(17) (a) Ziatdinov, V. R.; Oxgaard, J.; Mironov, O. A.; Young,
K. J. H.; Goddard, W. A., III; Periana, R. A. J. Am. Chem. Soc. 2006,
128, 7404–7405. (b) Davies, D. L.; Donald, S. M. A.; Macgregor, S. A.
J. Am. Chem. Soc. 2005, 127, 13754–13755. (c) Davies, D. L.; Donald,
S. M. A.; Al-Duaji, O.; Macgregor, S. A.; P€olleth, M. J. Am. Chem. Soc.
2006, 128, 4210–4211. (d) Bischof, S. M.; Ess, D. H.; Meier, S. K.;
Oxgaard, J.; Nielsen, R. J.; Bhalla, G.; Goddard, W. A., III; Periana, R. A.
Organometallics 2010, 29, 742–756.
14215
dx.doi.org/10.1021/ja2058154 |J. Am. Chem. Soc. 2011, 133, 14212–14215