A three-stage mechanistic model for ammonia-borane dehydrogenation by Shvo's catalyst

Article abstract of DOI:10.1021/om300562d

We propose a mechanistic model for three-stage dehydrogenation of ammonia-borane (AB) catalyzed by Shvo's cyclopentadienone-ligated ruthenium complex. We provide evidence for a plausible mechanism for catalyst deactivation and the transition from fast catalysis to slow catalysis and relate those findings to the invention of a second-generation catalyst that does not suffer from the same deactivation chemistry. The primary mechanism of catalyst deactivation is borazine-mediated hydroboration of the ruthenium species that is the active oxidant in the fast catalysis case. This transition is characterized by a change in the rate law for the reaction and changes in the apparent resting state of the catalyst. Also, in this slow catalysis situation, we see an additional intermediate in the sequence of boron, nitrogen species, aminodiborane. This occurs with concurrent generation of NH₃, which itself does not strongly affect the rate of AB dehydrogenation.

Full text of DOI:10.1021/om300562d

Article
pubs.acs.org/Organometallics
A Three-Stage Mechanistic Model for Ammonia−Borane
Dehydrogenation by Shvo’s Catalyst
Zhiyao Lu, Brian L. Conley, and Travis J. Williams*
Loker Hydrocarbon Research Institute and Department of Chemistry, University of Southern California, Los Angeles, California
90089-1661, United States
S
* Supporting Information
ABSTRACT: We propose a mechanistic model for three-stage
dehydrogenation of ammonia−borane (AB) catalyzed by Shvo’s
cyclopentadienone-ligated ruthenium complex. We provide
evidence for a plausible mechanism for catalyst deactivation and
the transition from fast catalysis to slow catalysis and relate those
findings to the invention of a second-generation catalyst that does
not suffer from the same deactivation chemistry. The primary
mechanism of catalyst deactivation is borazine-mediated hydro-
boration of the ruthenium species that is the active oxidant in the
fast catalysis case. This transition is characterized by a change in the rate law for the reaction and changes in the apparent resting
state of the catalyst. Also, in this slow catalysis situation, we see an additional intermediate in the sequence of boron, nitrogen
species, aminodiborane. This occurs with concurrent generation of NH3, which itself does not strongly affect the rate of AB
dehydrogenation.
INTRODUCTION
■
Hydrogen is an attractive alternative transportation fuel that has
appeal because it is carbon-free, is easily oxidized in fuel cells,
and is potentially available from water electrolysis.1 Although
pressurized hydrogen gas currently has some use in vehicles, its
practicality in cars is limited by fuel range, convenience, and
safety concerns.2 Particularly, hydrogen has low volumetric
energy density (5.6 MJ/L at 700 bar) despite high mass energy
density (120 MJ/kg for hydrogen).3 Thus, a highly weight-
efficient strategy to store hydrogen as condensed matter might
enable its translation more broadly into transportation
applications and consumer products.
Ammonia−borane (AB) is a promising material from which
to build a practical hydrogen storage system, because it has high
hydrogen density (19.6 wt %, ca. 9.9 MJ/L, 12.6 MJ/kg) and it
can release hydrogen under mild conditions (thermolysis,
hydrolysis, and catalysis; Figure 1).4 Transition-metal-catalyzed
dehydrogenation of ammonia−borane, particularly, is an area of
active research interest because it may enable a more efficient
fuel cycle than the well-known catalytic hydrolysis reaction that
forms ammonia, which is poisonous to fuel cells, and strong B−
O bonds, which are energetically costly to rereduce.5 Several
active transition-metal-based catalysts have been reported for
AB dehydrogenation reactions, in either heterogeneous or
homogeneous systems; these involve rhodium,6 iridium,7
ruthenium,8−10 nickel,11 palladium,12 and iron catalysts,13
among others14 (Figure 2).
Figure 1. Ammonia−borane dehydrogenation reactions and possible
products.
coordinative saturation of the reduced form of the Shvo system
would preclude coordination of aminoborane, NH2BH2, to the
catalyst, and thus disfavor the formation of insoluble oligomers,
[NH2BH2]n, which limits the hydrogen production of some
catalysts7,8 for ammonia−borane dehydrogenation to 1 equiv.16
Although this proposal seems to hold true,9 the catalyst begins
to deactivate after ca. 25% conversion in the first pass, which
renders this system irrelevant to practical implementation.
Our laboratory’s studies on catalytic AB dehydrogenation
have been focused on the reactivity of Shvo’s catalyst (12,
Scheme 1)15 and its relatives. By analogy to the established
mechanism for alcohol oxidation with 12, we presumed that the
Received: June 19, 2012
Published: August 30, 2012
© 2012 American Chemical Society
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Figure 2. Transition-metal catalysts for AB dehydrogenation.
generation system then enables access to high weight efficiency
dehydrogenation of ammonia−borane.10
Scheme 1. Dehydrogenation of AB with Shvo’s Catalyst, 12
RESULTS AND DISCUSSION
■
The kinetic profile of AB dehydrogenation with 12 shows
complicated behavior that we deconvoluted into a sequence of
three limiting cases (Figure 3).9 In our prior work we described
cases 1 and 2 in detail:9 catalyst initiation and fast catalysis,
respectively. The former is the case of low conversion and zero
[borazine] with ruthenium beginning in its dimeric form. This
can be easily studied in isolation, because these are the
conditions at the beginning of the reaction and because
initiation occurs quickly at 55 °C, where the catalysis is slow.
The second case is the one in which AB conversion and
[borazine] are low and the ruthenium in the system is no
longer in the form of its dimeric precursor. This case can easily
This text examines the molecular events that deactivate the
Shvo catalyst in ammonia−borane dehydrogenation. These
include hydroboration of the active, oxidizing form of the
catalyst by borazine, which involves the borylation of the
catalyst’s ligand oxygen atom so that the turnover-limiting H−
H bond-forming step is no longer accessible. Ultimately, we
addressed this problem by designing a second-generation
system, 13,17 that does not rely on an oxygen center as a proton
acceptor in the same way as the Shvo catalyst. This second-
Figure 3. (left) 11B NMR data showing consumption of AB in the presence of 2.5 mol % 12 in a sealed J. Young NMR tube. (right) Eudiometer data
showing production of hydrogen gas in the presence of 5.0 mol % 12 and 2.0 mol % ethanol in 2/1 diglyme/benzene-d6 at 70 °C.
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be studied in isolation by either (1) allowing the reaction to
incubate at room temperature until it is initiated, i.e. 12’s
characteristic μ-H peak (δ(1H) −17.7 ppm) is consumed, or
(2) delivering the ruthenium as dimer 18 (vide infra, Scheme
2). The third case, slow catalysis, occurs under the condition
that [borazine] ≈ [Ru atoms]. In this case the rate of
dehydrogenation catalysis drops precipitously and the catalyst
“dies”. This case can be generated in isolation by adding a
catalytic portion (1 equiv versus [Ruatom]) of borazine to the
reaction mixture at the outset of dehydrogenation.9
Figure 3 shows (left) AB consumption and (right) H2
generation as functions of time for trials of dehydrogenation
that cycle through all three of its mechanistic cases.9 This is
particularly evident from AB consumption: AB is consumed
slowly as the catalyst initiates (case 1) and then it is consumed
more quickly once dimer 12 is cleaved. Ultimately, [borazine]
increases and the catalysis slows down.
Understanding the mechanism of dehydrogenation in the
high [borazine] case and the catalyst deactivation process is
particularly important for the following reasons. (1) The kinetic
profile of the AB consumption in the high [borazine] case is
that observed when the catalyst is reused, and reusability is
essential for practical applications. (2) Borazine is an
unavoidable intermediate in AB dehydrogenation if ≥2 mol
equiv of hydrogen is released, yet the coordination chemistry of
borazine is not well studied in the context of catalytic AB
dehydrogenation systems.3 Understanding the mechanism of
this deactivation is important for the design of more efficient
catalytic systems and, ultimately, a broadly useful solution to
reversible hydrogen storage on ammonia−borane.
In this experiment we see that conversion of 18 to 16 reaches
completion within 5 min at 50 °C (see the Supporting
Information). This corresponds to a rate constant of >10−2 s−1
at 50 °C, which is faster than the rate of catalyst initiation (10−3
s−1 at 55 °C). Thus, the dissociation of dimer 12 is rate-limiting
in catalyst initiation. This result is in accordance with Shvo’s
considerable dissociation enthalpy: 28.8 kcal/mol in toluene in
the absence of AB.18b
II. Fast Catalysis (Case 2). After the catalyst initiation, the
kinetic profile of AB consumption displays fast, linear kinetics
through ca. 20−30% conversion. In these conditions, the
reaction has a zero-order dependence on [AB] and first-order
dependence on the catalyst’s [Ru], as determined by the
respective zero and unity slopes of plots of ln kobs versus ln
9
1
[AB] and ln [Ru]. Throughout this case, H NMR shows a
persistent monomeric ruthenium hydride at δ(1H) −10 ppm,
which is plausibly the resting state of the catalyst. This is
consistent with H−H bond formation as the turnover-limiting
step in fast catalysis. This assignment is consistent with the
observed kinetic dependencies of [AB]0 and [Ru]1. We further
observe first-order dependence on [EtOH], which is consistent
with the transition state model established by Casey for
stoichiometric hydrogen loss from 16.18 Thus, we adopt Casey
and Cui’s geometry for ethanol-mediated H−H bond formation
from 16 as the turnover-limiting transition state of this catalysis
(Scheme 3). In sum, the observed rate law in this case of the
reaction is −d[AB]/dt = kobs[Ru][EtOH].
Scheme 3. Proposed Mechanism of Fast Catalysis (Case 2)
I. Catalyst Initiation (Case 1). A short period in the
beginning of the reaction (ca. 2% conversion) is an initiation
period in which AB consumption is slow. This situation can be
studied in isolation by monitoring the reaction at a temperature
well below that needed for fast catalysis, 70 °C. Thus, upon
heating to 55 °C, the bridging hydride in 12 (δ(1H) −17.7
ppm) is replaced by the hydride of monomeric species 16
(δ(1H) −10.0 ppm) at a rate of 7.96(21) × 10−4 s−1.9 This
indicates that Shvo’s catalyst, 12, dissociates to its reduced
monomer 16 and (presumably) oxidized monomer 17
1
(Scheme 2A). H NMR integrations show that 2 equiv of 16
is formed with consumption of 12; thus, the reduction of 17 is
rapid relative to dissociation of 12.
The finding of rapid reduction of 17 by ammonia−borane is
supported by the relatively rapid rate of reduction of 18, a
stable dimer of 17, under analogous conditions (Scheme 2B).
III. Slow Catalysis (Case 3). Kinetics of Ammonia−
Borane Dehydrogenation in the Slow Catalysis Case. Onset
of the slow catalysis conditions, i.e. catalyst deactivation, is
characterized by the appearance of curvature in the time course
plot of [AB], and it becomes more likely as [borazine] rises.
The conditions of slow catalysis cause the emergence of
multiple κ1-Ru−H hydride peaks (from δ(1H) −9 to −10
ppm), which occurs simultaneously with exponential decay
behavior in [AB]. We believe that these correspond,
respectively, to (a) new resting state(s) of the catalyst and
ammonia−borane’s role in a new turnover-limiting step.
Understanding this mechanism is essential to our studies on
catalyst reuse and, we infer, spent fuel regeneration.
a
Scheme 2. Catalyst Initiation
An essential feature of the high [borazine] case of the
reaction is that the catalyst is reusable within the limits of its
kinetics. Thus, if a completed reaction mixture is treated with a
new aliquot of ammonia−borane, dehydrogenation will
recommence upon heating, and the reaction’s kinetic profile
a
A: scheme for catalyst initiation. B: rapid formation of 16 from 18.
[AB] is 0.42 M in benzene-d6 solution, and [Ru]2 is 5 mol % to AB.
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will follow slow catalysis behavior wherein the rate of hydrogen
production is too slow to be useful.9 It is easy to believe that, at
the conclusion of the reaction, the concentration of [17] is very
small, so that dimer 12 is not re-formed, and it is unnecessary
to repeat catalyst initiation (case 1) in catalyst reuse
experiments. However, the absence of fast catalysis must result
from a practically irreversible deactivation of the catalyst during
the transition from fast catalysis to slow catalysis in the first run.
The exponential decay shape of the kinetic profile of AB
consumption suggests that the reaction is now first order in
[AB] in the slow catalysis regime (Table 1, left), as verified by a
conundrum, however, which is that the product of the two
single-label KIEs should equal the double-label KIE; in this case
we have 1.46(3) × 1.07(5) = 1.56(6), which is well below the
observed value of 2.30(4). Along these lines, Casey’s group has
reported H/D exchange of the Ru−H group in Shvo’s catalyst
with D2 (or vice versa) of 16-Tol in THF without substituting
the corresponding ligand O−H.18 Similarly, in two parallel runs
of ND3BH3 dehydrogenation, similar portions of HD and H2
were formed. The presence of H2 (and by symmetry D2)
implies the availability of a mechanism for proton/hydride
exchange under our catalytic conditions. On the basis of our
observations and their result, we conducted an experiment of
ND3BD3 dehydrogenation with 1 atm of H2 gas applied to the
solution. HD was formed during this reaction (Scheme 4A),
Table 1. Ammonia−Borane Consumption as a Function of
a
[AB] and [Ru] in Slow Catalysis
Scheme 4. H/D Exchange Experiments
b
c
AB (M) rate (kobs) (s−1
)
amount of 12 (mol %) rate (kobs) (s−1
)
0.42
0.52
0.73
0.94
5.99(12) × 10−5
7.09(19) × 10−5
1.05(2) × 10−4
1.35(4) × 10−4
2.5
4.31(7) × 10−5
3.75
5.0
5.13(6) × 10−5
5.99(12) × 10−5
7.49(12) × 10−5
which necessitates an H/D crossover mechanism involving the
final product, H2. We believe that the mechanism of this is the
same as Casey’s H/D exchange mechanism, except we suggest
that this mechanism is available to the resting state(s) of our
catalyst. To test this latter hypothesis, we treated borylated
ruthenium complex 26 with 1 atm of D2 under conditions
analogous to our catalytic reactions. We observed the formation
of HD at room temperature in 5 min and complete deuteration
in the hydride position of 26 within 1 h at 60 °C (Scheme 4B).
This shows us that there is a mechanism for H/D exchange of
the ruthenium hydride in an O-borylated homologue of the
Shvo system. This result provides an explanation for the small
experimental kNHBH/kNHBD value and the mismatch between
our observed kNHBH/kNDBD and the value predicted by the
separate values for proton and hydride: because there is a facile
mechanism for H/D exchange, an isotopic kinetic resolution is
possible.
Mechanistic Proposal. We propose, on the basis of NMR
observations, that fast catalysis ends (i.e., catalyst deactivaiton
occurs) because borazine undergoes a hydroboration with
ruthenium intermediate 17 to give the deactivated complex 21,
which further converts to other derivatives (Scheme 5).9
Reactions analogous to the addition of 3 to 17 are known from
the Casey18a and Clark22 laboratories (Scheme 6). Casey has
shown hydrosilylation of the Shvo scaffold by triethylsilane.
This adduct, 24-Tol, has a δ(1H) value of −9.20 ppm in
benzene-d6. Similarly, Clark has shown hydroboration of the
Shvo complex with pinacol−and catechol−boranes in high
yield at mild temperature. These adducts have δ(1H) values of
−9.33 and −9.26 ppm in benzene-d6, respectively.
7.5
a
Data calculated from 11B NMR monitored kinetic studies at 70 °C.
[Ruatom]0 = 42.0 mM. [AB]0 = 0.42 M.
b
c
plot of ln kobs versus ln [AB] with a slope of 1.03(4) as
recorded in the case of high [borazine]. This draws a significant
contrast to fast catalysis, where such a plot was of zero slope,
and indicates that, unlike fast catalysis, once the catalyst
deactivates, the turnover-limiting step for further turnover
involves conversion of the resting species by 1 equiv of
ammonia−borane.
The kinetic order in ruthenium also changes upon the onset
of slow catalysis conditions. [Ru] is first order in the fast
catalysis case, but it becomes half order in the high [borazine]
case: rate constants were measured for a series of [Ru]
concentrations in AB dehydrogenation in the presence of
borazine, which gave a plot of ln kobs versus ln [Ru] with a slope
of 0.50(2) (Table 1, right). This is analogous to 12-catalyzed
alcohol oxidation,19 wherein apparent half-order dependence
on [Ru] is a result of equilibrium between 16 + 17 and dimer
12, for which a plot of ln kobs versus ln [Ruatom] has a slope of
0.40(6).20
Isotope Effects. Kinetic isotope effects were determined for
dehydrogenation of selectively deuterium-labeled ammonia−
borane isotopologues21 ND3BH3, NH3BD3, and ND3BD3 in
high [borazine] conditions. Comparison of measured rate
constants for parallel runs at 70 °C gave kinetic isotope effects
of kNHBH/kNDBH = 1.46(3), kNHBH/kNHBD = 1.07(5), and kNHBH
NDBD = 2.30(4) (see the Supporting Information). The KIE in
/
k
NH is suggestive of a catalyst reactivation involving
participation of the NH in its turnover-limiting step. This
might be akin to a protonation of the resting state of the
catalyst by an acidic NH proton. These data present a
We propose that this hydroborated species can dimerize to
form a O−B−O and Ru−H−Ru bridged dimer (31, Scheme 5)
akin to the parent Shvo complex and Clark’s [(μ-(cat)B-
(C4Ar4O)2)Ru2(CO)4(μ-H)] dimer, which accounts for the
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reuse experiments. Because of their self-reactive nature, we are
unable to isolate these complexes directly, but when a mixture
of these materials is collected and excess borazine is
quantitatively removed under reduced pressure, the resulting
material can be isolated through aqueous workup. This
treatment cleaves any borazine rings remaining in the
borazine−catalyst complex(es) and affords a ruthenium-
containing adduct, ammonia complex 23,23 which can be
isolated in 42% yield. This observation gives strong evidence
that the deactivated catalyst, the one present in the slow
catalysis case, is covalently bound to a borazine moiety, because
NH3 could not have been delivered in any other plausible way.
An Analogue of the Deactivated Catalyst. Because our
efforts to isolate and characterize our proposed deactivated
catalyst were frustrated by its reactivity, we set about to devise a
borazine analogue with which we could hydroborate 18 and
generate a stable surrogate of the catalyst of the slow catalysis
case. The premise of this design was our hypothesis that
multiple equivalents of 17 are hydroborated by 1 equiv of
Scheme 5. Proposed Borazine-Mediated Hydroboration
1
borazine to yield multiple κ1-Ru−H signals in the H NMR
Scheme 6. Hydroboration and Hydrosilylation of Tol-18
spectrum. Along these lines, we prepared diazaborane 29 and
treated it with dimer 18. The result was near-quantitative
formation of 12 under rigorously anhydrous conditions
(Scheme 6). We infer from this result that proposed dimer
27, if formed, apparently loses a borabenzoimidazole rapidly to
regenerate 12. In contrast, hydroboration of 18 with
catecholborane to form 26 is facile and gives an oxygen-
substituted analogue of our proposed deactivated catalyst that is
free of N−H groups.
In situ preparation of 26 in a benzene/diglyme solution
affords an opportunity to compare the rate and kinetic profile
of 26-catalyzed ammonia−borane dehydrogenation with those
of the slow catalysis case (Scheme 7). The kinetic profiles each
Scheme 7. Synthesis of a Mechanistic Analogue for the
Proposed Deactivated Catalyst Complex
observed half-order kinetic dependence on [Ru]. These species
can re-enter the catalytic cycle if the B−O bond affixing
borazine to the catalyst is cleaved in the presence of ammonia−
borane, which accounts for its first-order kinetic dependence.
We do not know the mechanism of ammonia−borane’s
involvement in this step.
Catalyst Deactivation. We conducted a series of experi-
ments directly to interrogate our proposal for the mechanism of
slow catalysis, yet we observe that the proposed complex 21 is
not stable to isolation. Borazine was added to dimer 18 at room
temperature, and a bridging hydride peak formed at the
beginning of the reaction (δ(1H) −18.3) was then consumed in
2 min (Scheme 5). This was replaced by a set of 13 hydride
peaks from δ(1H) −9.3 to −10.0 ppm, which correspond to κ1-
Ru−H groups such as 22. We suspect that these signals
correspond to multiple hydroboration events on a single
borazine or ring-opened borazine derivatives.
We can create slow catalysis case conditions at the beginning
of a dehydrogenation reaction very simply by adding 1 mol
equiv of borazine relative to [Ruatom] to the reaction mixture
prior to heating. Under these conditions the kinetic profile of
the reaction does not show any properties of initiation or fast
catalysis but proceeds directly to the rate and rate law of slow
catalysis.9 This is strong evidence indicating that borazine is the
agent that causes catalyst deactivation and aptly accounts for
the instant slow catalysis situation that is observed in catalyst
appear first order in AB, but the rate for dehydrogenation with
catalyst precursor 26 is faster than slow catalysis by a factor of
ca. 3-fold. This faster rate could be a result of a more labile O−
B bond between the catalyst’s hydroxycyclopentadiene and the
corresponding boranes. A plot of ln kobs versus ln [Ru] gave a
slope of 0.51(3) (Table 2), which is in agreement with the
measured [Ru] dependence for slow catalysis. These data show
us that the dimerization behavior that we see in the slow
catalysis case is effectively recreated in borylated analogue 26.
Taken together, these data provide good anecdotal evidence
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Table 2. Ammonia−Borane Dehydrogenation Catalyzed by
Borylated Complex 26
Table 3. Rate of Slow Catalysis in the Presence and Absence
of BH3
a
conditions
BH3·THF
slow catalysis kobs (s−1
)
amount of 26 (mol %)
rate (s−1
)
1.57(7) × 10−4
7.0(13) × 10−5
6.3(6) × 10−5
5.0
7.5
10
1.26(3) × 10−4
1.50(2) × 10−4
1.82(2) × 10−4
2.18(2) × 10−4
THF only
b
parent conditions
a
Data calculated from 11B NMR kinetic studies at 70 °C. Smoothed
curves are empirical fits; kobs values shown are for slow catalysis, not
the entire curve. See the Supporting Information. A parallel run under
parent conditions has kobs in statistical agreement with others reported
15
b
that the deactivated catalyst is an O-borylated form of the
precatalyst.
herein.
Other Reaction Products and Intermediates in the Slow
Catalysis Case. Ammonia−borane dehydrogenation catalyzed
by 12 generates multiple B, N intermediates throughout the
reaction. In fast catalysis the intermediates detected by 11B
NMR are the same as those observed for other catalysts that are
known to liberate multiple equivalents of hydrogen from
ammonia−borane:9−11 AB → branched cyclotriborazane (2) →
borazine (3) (Figure 1A). In the slow catalysis case, however, a
new species appears, aminodiborane 6 (Scheme 8). This
hydroborate and deactivate the catalyst in the same way as
borazine or catecholborane.
If free BH3 from the dissociation of ammonia−borane is
impacting the course of the reaction in the slow catalysis case,
then free NH3 must also be present, and NH3 is known to
modulate the reactivity of the Shvo system.23 Thus, we propose
a second mechanism of catalyst deactivation, which is reversible
formation of 23 by NH3 ligation to the catalyst.
To interrogate directly the reactivity of ammonia adduct 23,
we prepared it independently through the addition of ammonia
gas to 18 (Scheme 9). Analogous to the case for 12, the 11B
Scheme 8. Formation of 6
Scheme 9. Synthesis of Ammonia Adduct 23
species should be dehydrogenated to borazine, and in fact, in
reactions in which this species is an intermediate, borazine
remains the only product upon completion of the reaction. We
believe that 6 is an adduct formed from BH3, dissociated from
ammonia−borane, and NH2BH2, generated transiently after the
first dehydrogenation of ammonia−borane.16,24 We propose
that formation of 6 is reversible, and this is only a mechanistic
cul de sac, rather than an in-line intermediate in the
dehydrogenation sequence. To test this hypothesis, we added
(a) 0.5 equiv of 1 M BH3·THF and (b) a comparable volume of
THF to two otherwise identical runs of ammonia−borane
dehydrogenation with 12 (Table 3). A strong signal for 6 was
observed by 11B NMR in tube a in the beginning of the
reaction, much earlier than the first emergence of 6’s peak in
the THF control experiment (tube b). Both reactions
proceeded through fast catalysis at about the same rate (Figure
3). Furthermore, the rates of AB consumption in slow catalysis
case are similar, ca. 25% difference, which shows that although
there is a large excess of 6 in tube a in comparison to the
amount in tube b, this has a disproportionately small effect on
the rate of AB consumption. We therefore know that 6 goes on
to dehydrogenate to borazine and does not significantly
interfere with the rates of the steps in slow catalysis as it
forms and disappears. It further appears that BH3 does not
NMR kinetic profile of AB dehydrogenation with 23 appears to
have two distinct kinetic cases, one linear case, with a reaction
rate of 5.12 × 10−5 M s−1, and one exponential decay case, with
a rate constant of 4.22 × 10−4 s−1 (Table 4). This is similar to
the second and third cases (fast and slow catalysis) of
dehydrogenation with 12, but since 23 is monomeric, it stands
to reason that there should not be an initiation delay analogous
to that observed in reactions featuring 12. We account for this
behavior by proposing that NH3 reversibly can ligate 17 as
previously documented23 and thereby temporarily sequester it
from its catalytic roles. Thus, NH3 ligation provides a second
mechanism for catalyst deactivation, although this one appears
to be less deleterious than hydroboration of 17.
Homogeneous versus Heterogeneous Catalysis. We
propose that this reaction is homogeneous throughout its
duration on the basis on four observations. First, the reactor
maintains its homogeneous appearance through the duration of
the reactions. No metallic residue is observed. Second, the rate
of catalysis is not impacted by the addition of Hg(0). In
contrast, a mercury drop does inhibit the catalytic hydro-
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a
Table 4. [AB] Dehydrogenation with 12 and 23
fast catalysis
cat.
slow catalysis
cat.
rate (M s−1
)
kobs (s−1
)
12
23
1.47(9) × 10−4
5.12(3) × 10−5
12
23
3.06(39) × 10−4
4.23(33) × 10−4
a
Data calculated from 11B NMR kinetic studies. A 0.25 mol portion of
AB and 0.035 mol of [Ruatom] were added to 0.6 mL to diglyme/
benzene-d6. Both reactions were run at 70 °C. Black circles and
diamonds are kinetic profiles of reactions catalyzed by 23 and 12,
respectively.
Figure 4. [AB] dehydrogenation with 12 in 1,10-phenanthroline. Data
are calculated from 11B NMR kinetic studies. A 0.25 mol amount of
AB and 5 mol % of 12 were added to 0.6 mL of diglyme/benzene-d6 at
70 °C.
genation of benzene based on the [Ru3(μ2-H)3(η6-C6H6)(η6-
C6Me6)2(μ3-O)]+ catalyst precursor, which is part of the
evidence for heterogeneous reduction in that system.25 This
result is germane to the present discussion because it shows a
documented system wherein the mercury drop experiment was
effective with ruthenium. Still, the best evidence we have (third
point) for homogeneous catalysis remains the foregoing
kinetics data, in which the data remain pseudo first order in
the slow catalysis case through >90% conversion (>3 half-lives).
A fourth piece of evidence favoring homogeneous catalysis
comes from a quantitative poisoning experiment wherein the
reaction is run in the presence of a small portion of 1,10-
phenanthroline.
an analogous Rh4Cp*2.4Cl4Hc cluster has been shown to be the
likely active catalyst in the hydrogen of benzene with
[Cp*RhCl2]2, and this Rh4 cluster is deactivated by Hg(0).27
Further, that Rh4-based system is deactivated by 4 equiv of
1,10-phenanthroline (i.e., 1/1 phen/Rhatom), while ours is not.
A Second-Generation Catalyst. To avoid the problem of
catalyst O-borylation that is intrinsic to the deactivation of the
Shvo catalyst, we devised the second-generation system
highlighted in Scheme 10.10,17 One design feature of this
Quantitative poisoning is an experiment in which less than 1
molar equiv (relative to the proposed monomeric catalyst) is
introduced into the reaction, and one monitors the rate to see if
it is affected proportionally to the concentration of the poison.
If the drop in rate upon poisoning is disproportionally large,
this can be evidence for heterogeneous catalysis, because
<100% metal atoms (and often ≤50%)26 are on the surface of a
nanoparticle and thus ≤50% are available to be poisoned.
In this present case, two quantitative poisoning experiments
were conducted wherein 0.1 and 0.5 equiv of 1,10-phenanthro-
line (phen) relative to [Ruatom] were added to two otherwise
standard ammonia−borane dehydrogenation runs with catalyst
12 (0.25 mol of AB, 70 °C, diglyme/benzene-d6). Although
these are not first-order reactions, generally we see that phen
accelerates the reaction, apparently by prolonging the fast
catalysis portion of the reaction (Figure 4). In contrast, 0.5
equiv of phen (relative to ruthenium atoms) completely
quenches catalytic heterogeneous hydrogenation of benzene
based on the [Ru3(μ2-H)3(η6-C6H6)(η6-C6Me6)2(μ3-O)]+
catalyst precursor.25 Thus, we take our evidence to argue
against the formation of a ruthenium nanoparticle. The origin
of the acceleration behavior seems as if the phenanthroline
“poison” is protecting the catalyst from borazine-mediated
deactivation. This concept is further discussed in the
Supporting Information.
Scheme 10. A Second-Generation Catalyst
motif is that if the catalyst is borylated by analogy to 21, the
resulting borylated oxygen atom can be lost from the catalyst,
so that the catalytic species is not deactivated. Although we do
not know the mechanism of ammonia−borane dehydrogen-
ation with 13, we observe that it can liberate >2 equiv of H2
from ammonia−borane, even on repeated uses and with
exposure to the atmosphere, and it shows first-order
consumption of ammonia−borane throughout conversion.
A tetranuclear catalyst (e.g., Ru4Ln) is unlikely but cannot be
rigorously eliminated. Such a hypothesis is disfavored because
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reactive as their mechanistic predecessor in ammonia−borane
dehydrogenation. Presumably, these deactivated ruthenium
species are reactivated by ammonia−borane itself and proceed
to further ammonia−borane dehydrogenation. This observation
gives us insight into the higher reactivity of our second-
generation catalyst, 13.
CONCLUSION
In conclusion, we were able to propose a full picture of the
catalyst deactivation mechanism (Scheme 11). A full scheme
■
Scheme 11. Mechanitic Proposal of Slow Catalysis
EXPERIMENTAL SECTION
■
I. General Procedures. All air- and water-sensitive procedures
were carried out either in a Vacuum Atmospheres glovebox under
nitrogen (0.5−10 ppm of O2 for all manipulations) or using standard
Schlenk techniques under nitrogen. Deuterated NMR solvents were
purchased from Cambridge Isotopes Laboratories. Benzene-d6 and
diethylene glycol dimethyl ether (diglyme, J. T. Baker) were dried over
sodium benzophenone ketyl and distilled prior to use. Shvo’s catalyst
was purchased from Strem Chemicals. Ammonia−borane (NH3BH3,
AB) was purchased from Sigma Aldrich. Borazine was synthesized and
purified by the method used by Wideman and Sneddon.28 1H and 11
B
NMR spectra were obtained on a Varian 600 spectrometer (600 MHz
in 1H, 192 MHz in 11B) with chemical shifts reported in units of ppm.
All 1H chemical shifts are referenced to the residual 1H solvent
(relative to TMS). All 11B chemical shifts are referenced to BF3·OEt2
in diglyme in a coaxial external standard (0 ppm). NMR spectra were
taken in 8 in. J. Young tubes (Wilmad) with Teflon valve plugs. The
NMR tubes were shaken vigorously for several minutes with
chlorotrimethylsilane and then dried in vacuo on a Schlenk line
prior to use.
Caution! Extreme caution should be used when carrying out these
reactions, as the release of hydrogen can lead to sudden pressurization
of reaction vessels.
1
II. Mechanistic Studies Utilizing 11B and H NMR Spectros-
copy. In a typical reaction, 7.7 mg of AB was combined with Shvo’s
catalyst (12, 13.6 mg, 5 mol %) in a J. Young NMR tube while in a
glovebox under nitrogen. The AB and catalyst concentrations may be
varied. Diglyme (0.4 mL) and benzene-d6 (0.2 mL) were added to the
tube, as was the BF3 insert. The sample tube was immediately inserted
into a preheated NMR (70 °C), and the kinetic monitoring
commenced after quickly locking and shimming. Disappearance of
AB in the solution was monitored by the relative integration of its
characteristic peak in the 11B spectrum (−22 ppm) and the BF3·OEt2
standard. All spectra were processed using VNMRJ (version 2.3). The
acquisition involved a 1.67 s pulse sequence in which 4096 complex
points were recorded, followed by 1 s relaxation delay. To eliminate
B−O peaks from the borosilicate NMR tube and probe, the 11B FIDs
were processed with back linear prediction, ca. 5−15 points.
A. Determination of Catalyst Order in Conversion of AB in Case 3
(Slow Catalysis). The rate values for the slow catalysis case were
determined using 11B NMR in sealed NMR tubes, as described above.
Data treatments are shown in the Supporting Information. The
amount of AB was 7.7 mg (0.25 mmol), and catalyst concentrations
were varied (6.8, 10.2, 13.6, and 20.3 mg of 12, (2.5, 3.75, 5.0, and 7.5
mol %)). The results were plotted as a ln/ln relationship to determine
the order in catalyst (Table 1, right).
showing the origin of the slow catalysis case is illustrated in a
diagram in the Supporting Information. The upper-right half of
Scheme 11 is the fast catalysis mechanism, which is sketched in
Scheme 3. We propose that catalyst deactivation is caused by
the introduction of borazine, the product of selective
dehydrogenation, into solution. When the rate of reduction
of oxidized catalyst 17 by ammonia−borane becomes
competitive with its hydroboration, AB consumption ceases
to be linear. When hydroboration of 17 becomes fast relative to
reduction of 17, then 21 becomes a resting state of [Ru], and
the catalysis “dies”, as sketched in the lower-left half of the cycle
in Scheme 11. In this phase of the reaction, conversion has first-
order dependence on [AB], apparently because ammonia−
borane is needed to convert borazine species 21 back into an
active species.
Reversible ammonia ligation of 17 is almost certainly
happening throughout the reaction. We know, however, from
the rates of (a) fast catalysis and (b) slow catalysis with
ammonia complex 23 that 23 is a minor contributor to the total
[Ru atoms] in fast catalysis because (i) the addition of 1 equiv
of NH3 relative to ruthenium slows fast catalysis and (ii) the
rate of AB consumption is constant throughout the fast catalysis
case. We also know that NH3 ligation to ruthenium does not
significantly alter the rate of slow catalysis.
B. Determination of Order in AB in Case 3 (Slow Catalysis). The
rate values for the slow catalysis case were again determined using 11B
NMR in sealed NMR tubes. Data treatments are shown in the
Supporting Information. The amount of 12 was 13.6 mg (0.013
mmol), and AB concentrations were varied (7.7, 8.7, 13.5, and 17.4 mg
AB (0.42, 0.53, 0.73, and 0.94 M)). The results were plotted as a ln/ln
relationship to determine the order in ammonia-borane (Table 1, left).
C. Kinetic Isotope Effects in Case 3 (Slow Catalysis). To determine
the kinetic isotope effects on the reaction rate, 8.5 mg of ND3BH3, 8.5
mg of NH3BD3, or 9.2 mg ND3BD3 (0.25 mmol) was added to a J.
Young NMR tube. To each was added 13.6 mg of 12 (5 mol %),
diglyme (0.4 mL), and benzene-d6 (0.2 mL). Again, we analyzed [AB]
vs time for the third case of the reaction conditions. Data treatments
are shown in the Supporting Information. KIEs were determined from
the quotient of protic AB kobs (5.99(12) × 10−5 s−1) divided by the
In summary, the mechanism of AB dehydrogenation
catalyzed by Shvo catalyst 12 was investigated. This reaction
initiates with dissociation of the dimeric precatalyst 12 and then
goes through a fast dehydrogenation reaction wherein the H−
H bond formation is the turnover-limiting step. As the
concentration of borazine increases, it adds to the reactive
form of the catalyst to give ruthenium species, which are not as
1
deuterated AB kobs. The HD signal found in the H NMR resulted
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from H-D exchange and is also shown in Figure S5 (Supporting
Information).
ACKNOWLEDGMENTS
■
We thank Richard Finke and Ercan Bayram (Colorado State)
for insightful discussions. This work was sponsored by the
National Science Foundation (CHE-1054910) and the Hydro-
carbon Research Foundation. We are grateful to the National
Science Foundation (DBI-0821671, CHE-0840366), the Na-
tional Institutes of Health (1 S10 RR25432), and the University
of Southern California for their sponsorship of NMR
spectrometers at USC.
D. Kinetics for AB Dehydrogenation in the Presence of Added
BH3·THF. To examine the role of aminodiborane (6) in the mechanism
of AB dehydrogenation, we manipulated its concentration by adding
(tube A) 0.125 mL of 1 M BH3·THF (0.125 mmol, 0.5 equiv to AB)
or (tube B) 0.125 mL of THF to otherwise typical AB dehydrogen-
ation reactions (7.7 mg (0.25 mmol) of AB, 13.6 mg (5 mol %) of 12,
0.6 mL of 2/1 diglyme/benzene-d6). 11B NMR showed 6 (−26.7 ppm)
in tube A at the beginning of the reaction (see the Supporting
Information). The kobs values for these runs were determined using 11
B
NMR; data treatment is shown in the Supporting Information.
E. Kinetics for AB Dehydrogenation by NH3-Ligated Species 23.
Rate values for 23-catalyzed dehydrogenation run at 70 °C were
determined using 11B NMR as shown in the Supporting Information.
In this reaction 7.7 mg of AB (0.25 mmol) was combined with (tube
A) 20.7 mg of 23 (35 μmol, 14 mol %) and (tube B) 38.0 mg of 12
(35 μmol, 14 mol %). Data are shown in Table 4.
F. 1,10-Phenanthroline Poisoning Experiments. Fractional poison-
ing experiments employing separately 0.5, 0.1, and 0 mol equiv of 1,10-
phenanthroline relative to [Ruatom] were done under conditions
otherwise identical with our standard conditions for ammonia−borane
dehydrogenation by 12.
1,10-Phenanthroline (2.2 mg, 12 μmol, 50 mol % versus [Ruatom])
was added to a solution of 7.7 mg of ammonia−borane (0.25 mmol)
and 13.6 mg of 12 (25 μmol, 5 mol % versus AB) in 0.6 mL of 2/1
diglyme/benzene-d6. The rates for both the fast and slow catalysis
cases at 70 °C were determined using 11B NMR, as shown in the
Supporting Information.
1,10-Phenanthroline (2.2 mg) was dissolved in 0.5 mL of diglyme to
make a 1,10-phenanthroline stock solution. A portion of this solution
(0.1 mL, 12 μmol, 10 mol % 1,10-phenanthroline versus [Ruatom]) was
added to a solution of 7.7 mg of ammonia−borane (0.25 mmol) and
13.6 mg of 12 (25 μmol, 5 mol % versus AB) in 0.5 mL of 2/1
diglyme/benzene-d6. The rates for both the fast and slow catalysis
cases at 70 °C were determined using 11B NMR, as shown in the
Supporting Information.
A further fractional poisoning regarding the role of 1,10-
phenanthroline in the 26-catalyzed dehydrogenation of ammonia−
borane was also conducted. These data show little change in reactivity
upon addition of the poison (see the Supporting Information).
III. Preparative and Spectroscopic Details for Adduct 23. The
Ru−NH3 adduct 23 was prepared by delivering ammonia gas to a
benzene (5 mL) solution of 18 (50 mg, 0.046 mmol). The reaction
mixture was stirred at room temperature for 15 min. A black
precipitate was filtered out and successively washed with deionized
water, acetone, benzene, and hexanes. The solid was then dried under
vacuum to give a pale gray powder in 59% yield (30 mg).
REFERENCES
■
(1) (a) Imarisio, G. Int. J. Hydrogen Energy 1981, 6, 153−158.
(b) Zoulias, E.; Varkaraki, E.; Lymberopoulos, N.; Christodoulou, C.
N.; Karagiorgis, G. N. TCJST 2004, 4, 41−71. (c) Liu, R.-S.; Zhang, L.;
Sun, X.; Liu, H.; Zhang, J. Water Electrolysis for Hydrogen
Generation, In Electrochemical Technologies for Energy Storage and
Conversion; Millet, P., Ed., Wiley-VCH: Weinheim, Germany, 2011;
Vol. 1, p 2.
(2) United States Department of Energy, Hydrogen Fuel Cells and
Infrastructure Technology Program. Multi-Year Research Development
and Demonstration Plan, General Printing Office: Washington, DC,
2009; Section 3.3.
(3) Satyapal, S.; Petrovic, J.; Read, C.; Thomas, G.; Ordaz, G. Catal.
Today 2007, 120, 246−256.
(4) For a review, see: (a) Staubitz, A.; Robertson, A. P. M.; Manners,
I. Chem. Rev. 2010, 110, 4079−4124. For thermolysis of ammonia−
borane see: (b) Baitalow, F.; Wolf, G.; Jaeicke-Rossler, K.; Leitner, G.
Thermochim. Acta 2006, 445, 121−125. (c) Baitalow, F.; Baumann, J.;
Wolf, G.; Jaenicke-Rossler, K.; Leitner, G. Thermochim. Acta 2002,
391, 159−168. (d) Wolf, G.; Baumann, J.; Baitalow, F.; Hoffmann, F.
P. Thermochim. Acta 2000, 343, 19−25. (e) Wang, J. S.; Geanangel, R.
A. Inorg. Chim. Acta 1988, 148, 185−190. (f) Bluhm, M. E.; Bradley,
M. G.; Butterick, R., III; Kusari, U.; Sneddon, L. G. J. Am. Chem. Soc.
2006, 128, 7748−7749. (g) Rassat, S. D.; Aardahl, C. L.; Autrey, T.;
Smith, R. S. Energy Fuels 2010, 24, 2596−2606. (h) Nylen, J.; Sato, T.;
Soignard, E.; Yarger, J. L.; Stoyanov, E.; Haussermann, U. J. Chem.
Phys. 2009, 131, 104506. (i) Heldebrant, D. J.; Karkamakar, A.; Hess,
N. J.; Bowden, M.; Rassat, S.; Zheng, F.; Rappe, K.; Autrey, T. Chem.
Mater. 2008, 20, 5332−5336. For hydrolysis, see: (j) Marder, T. B.
Angew. Chem., Int. Ed. 2007, 46, 8116−8118. (k) Yan, J. M.; Zhang, X.
B.; Akita, T.; Haruta, M.; Xu, Q. J. Am. Chem. Soc. 2010, 132, 5326−
5327. (l) Jiang, H. L.; Umegaki, T.; Akita, T.; Zhang, X. B.; Haruta, M.;
Xu, Q. Chem. Eur. J. 2010, 16, 3132−3137. (m) Ramachandran, P. V.;
Gagare, P. D. Inorg. Chem. 2007, 46, 7810−7817.
1H NMR (pyridine-d5, 600 MHz): δ 8.05 (d, JHH = 7.1 Hz, 4H, Ph),
7.47 (d, JHH = 7.1 Hz, 4H, Ph), 7.18 (t, JHH = 7.1 Hz, 4H, Ph), 7.13
(m, 6H, Ph), 7.07 (t, JHH = 7.1 Hz, 4H, Ph), 4.25 (br. s, 3H, NH3).
13C{1H} NMR (pyridine-d5, 150 MHz): δ 202.5 (CO), 165.5 (C1 of
Cp), 134.9 (Ph), 133.4 (Ph), 133.2 (Ph), 131.3 (Ph), 128.5 (Ph),
128.4 (Ph), 128.3 (Ph), 126.9 (Ph), 104.3 (C2,5 of Cp), 83.0 (C3,4 of
Cp). Data are consistent with a known compound.23
(5) (a) Stephens, F. H.; Pons, V.; Baker, R. T. Dalton Trans. 2007,
2613−2626. (b) Smythe, N. C.; Gordon, J. C. Eur. J. Inorg. Chem.
2010, 509−521. (c) Hamilton, C. W.; Baker, R. T.; Staubitz, A.;
Manners, I. Chem. Soc. Rev. 2009, 38, 279−293. (d) Sutton, A. D.;
Burrell, A. K.; Dixon, D. A.; Gardner., E. B.; Gordon, J. C.; Nakagawa,
T.; Ott, K. C.; Robinson, J. P.; Vasiliu, M. Science 2011, 331, 1426−
1429.
(6) (a) Jaska, C. A.; Templ, 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) Sgrestha, R. P.; Diyabalanage, H. V.
K.; Semelsberger, T. A.; Ott, K. C.; Burrell, A. K. Int. J. Hydrogen
Energy 2009, 34, 2616−2621. Regarding borane−Rh coordination,
see: (d) Douglas, T. M.; Chaplin, A. B.; Weller, A. S. J. Am. Chem. Soc.
2008, 130, 14432−14433. (e) Alcaraz, G.; Sabo-Etienne, S. Angew.
Chem., Int. Ed. 2010, 49, 7170−7179.
ASSOCIATED CONTENT
■
S
* Supporting Information
Text and figures giving further experimental details and data.
This material is available free of charge via the Internet at
http://pubs.acs.org.
(7) Denny, M. C.; Pons, V.; Hebdon, T. J.; Heinekey, M.; Goldberg,
K. I. J. Am. Chem. Soc. 2006, 128, 12048−12049.
AUTHOR INFORMATION
■
(8) Blaquiere, N.; Diallo-Garcia, S.; Gorelsky, I.; Black, A.; Fagnou, K.
Corresponding Author
*E-mail: travisw@usc.edu.
J. Am. Chem. Soc. 2008, 130, 14034−14035. (f) Kaβ, M.; Fridrich, A.;
̈
Drees, M.; Schneider, S. Angew. Chem., Int. Ed. 2009, 48, 905−907.
(9) Conley, B. L.; Williams, T. J. Chem. Commun. 2010, 46, 4815−
4817.
Notes
The authors declare no competing financial interest.
6713
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Organometallics
Article
(10) Conley, B. L.; Guess, D.; Williams, T. J. J. Am. Chem. Soc. 2011,
133, 14212−14215.
(11) Keaton, R. J.; Blacquiere, J. M.; Baker, R. T. J. Am. Chem. Soc.
2007, 129, 1844−1845.
(12) Kim, S.-K.; Han, W.-S.; Kim, T.-J.; Kim, T.-Y.; Nam, S. W.;
́
Mitoraj, M.; Piecos, Ł.; Michalak, A.; Hwang, S.-J.; Kang, S. O. J. Am.
Chem. Soc. 2010, 132, 9954−9955.
(13) (a) Vance, J. R.; Robertson, A. P. M.; Lee, K.; Manners, I. Chem.
Eur. J. 2011, 17, 4099−4103. (b) Baker, R. T.; Gordon, J. C.;
Hamilton, C. W.; Henson, N. J.; Lin, P. H.; Maguire, S.; Murugesu,
M.; Scott, B. L.; Smythe, N. C. J. Am. Chem. Soc. 2012, 134, 5598−
5609.
(14) (a) Chapman, A. M.; Haddow, M. F.; Wass, D. F. J. Am. Chem.
Soc. 2011, 133, 8826−8829.
(15) Conley, B. L.; Pennington-Boggio, M. K.; Boz, E.; Williams, T. J.
Chem. Rev. 2010, 110, 2294−2312.
(16) (a) Wright, W. R. H.; Berkeley, E. R.; Alden, L. R.; Baker, R. T.;
Sneddon, L. G. Chem. Commun. 2011, 47, 3177−3179. (b) Alcaraz, G.;
Vendier, L.; Clot, E.; Sabo-Etienne, S. Angew. Chem., Int. Ed. 2010, 49,
918−920.
(17) Conley, B. L.; Williams, T. J. J. Am. Chem. Soc. 2010, 132,
1764−1765.
(18) (a) Casey, C. P.; Singer, S. W.; Powell, D. R.; Hayashi, R. K.;
Kavana, M. J. Am. Chem. Soc. 2001, 123, 1090−1100. (b) Casey, C. P.;
Beetner, S. E.; Johnson, J. B. J. Am. Chem. Soc. 2008, 130, 2285−2295.
(c) Casey, C. P.; Johnson, J. B.; Singer, S. W.; Cui, Q. J. Am. Chem. Soc.
2005, 127, 3100−3109.
(19) Johnson, J. B.; Backvall, J.-E. J. Org. Chem. 2003, 68, 7681−
̈
7684.
(20) Thorson, M. K.; Klinkel, K. L.; Wang, J.; Williams, T. J. Eur. J.
Inorg. Chem. 2009, 295−302.
(21) Hu, M. G.; Van Paasschen, J. M.; Geanangel, R. A. J. Inorg. Nucl.
Chem. 1997, 39, 1247−1250.
(22) Koren-Selfridge, L.; Query, I. P.; Hanson, J. A.; Isley, N. A.;
Guzei, I. A.; Clark, T. B. Organometallics 2010, 29, 3896−3900.
(23) Hollmann, D.; Jiao, H.; Spannenberg, A.; Bahn, S.; Tillack, A.;
̈
Parton, R.; Altink, R.; Beller, M. Organometallics 2009, 28, 473−479.
(24) Stephens, F. H.; Pons, V.; Baker, R. T. Dalton Trans. 2007,
2613−2626.
(25) Hagen, C. M.; Vieille-Petit, L.; Laurenczy, G.; Suss-Fink, G.;
̈
Finke, R. G. Organometallics 2005, 24, 1819−1831.
(26) (a) Widegren, J. A.; Finkie, R. G. J. Mol. Catal. 2003, 198, 317−
341. (b) Kovac
Paris, Chim. 2000, 3, 601.
́ ́ ́
s, G.; Nadasdi, L.; Joo, F.; Laurenczy, G. C. R. Acad. Sci.
(27) Bayram, E.; Linehan, J. C.; Fulton, J. L.; Roberts, J. A. S.;
̈
Szymczak, N. K.; Smurthwaite, T. D.; Ozkar, S.; Balasubramanian, M.;
Finke, R. G. J. Am. Chem. Soc. 2011, 133, 18889−18902.
(28) Wildeman, T.; Sneddon, L. G. Inorg. Chem. 1995, 34, 1002−
1003.
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