CO Displacement in an Oxidative Addition of Primary Silanes to Rhodium(I)

Article abstract of DOI:10.1021/acs.inorgchem.8b03425

The rhodium dicarbonyl {PhB(Ox Me₂)₂Im^Mes}Rh(CO)₂ (1) and primary silanes react by oxidative addition of a nonpolar Si-H bond and, uniquely, a thermal dissociation of CO. These reactions are reversible, and kinetic measurements model the approach to equilibrium. Thus, 1 and RSiH₃ react by oxidative addition at room temperature in the dark, even in CO-Saturated solutions. The oxidative addition reaction is first-Order in both 1 and RSiH₃, with rate constants for oxidative addition of PhSiH₃ and PhSiD₃ revealing k_H/k_D a 1. The reverse reaction, reductive elimination of Si-H from {PhB(Ox Me₂)₂Im^Mes}RhH(SiH₂R)CO (2), is also first-Order in [2] and depends on [CO]. The equilibrium concentrations, determined over a 30 °C temperature range, provide ?"H° = a'5.5 ± 0.2 kcal/mol and ?"S° = a'16 ± 1 cal·mol^-1K^-1 (for 1 a?., 2). The rate laws and activation parameters for oxidative addition (?"H^a§§ = 11 ± 1 kcal·mol^-1 and ?"S^a§§ = a'26 ± 3 cal·mol^-1·K^-1) and reductive elimination (?"H^a§§ = 17 ± 1 kcal·mol^-1 and ?"S^a§§ = a'10 ± 3 cal·mol^-1K^-1), particularly the negative activation entropy for both forward and reverse reactions, suggest the transition state of the rate-Determining step contains {PhB(Ox Me₂)₂Im^Mes}Rh(CO)₂ and RSiH₃. Comparison of a series of primary silanes reveals that oxidative addition of arylsilanes is ca. 5× faster than alkylsilanes, whereas reductive elimination of Rh-Si/Rh-H from alkylsilyl and arylsilyl rhodium(III) occurs with similar rate constants. Thus, the equilibrium constant K_e for oxidative addition of arylsilanes is >1, whereas reductive elimination is favored for alkylsilanes.

Full text of DOI:10.1021/acs.inorgchem.8b03425

Article
pubs.acs.org/IC
Cite This: Inorg. Chem. XXXX, XXX, XXX−XXX
CO Displacement in an Oxidative Addition of Primary Silanes to
Rhodium(I)
Abhranil Biswas,†,‡ Arkady Ellern,‡ and Aaron D. Sadow*,†,‡
†Department of Chemistry, Iowa State University, 1605 Gilman Hall, 2415 Osborn Drive, Ames, Iowa 50011, United States
‡United States Department of Energy Ames Laboratory, Iowa State University, Ames, Iowa 50011, United States
S
* Supporting Information
ABSTRACT: The rhodium dicarbonyl {PhB(OxMe )2ImMes}Rh-
2
(CO)2 (1) and primary silanes react by oxidative addition of a
nonpolar Si−H bond and, uniquely, a thermal dissociation of CO.
These reactions are reversible, and kinetic measurements model
the approach to equilibrium. Thus, 1 and RSiH3 react by oxidative
addition at room temperature in the dark, even in CO-saturated
solutions. The oxidative addition reaction is first-order in both 1
and RSiH3, with rate constants for oxidative addition of PhSiH3
and PhSiD3 revealing kH/kD ∼ 1. The reverse reaction, reductive
elimination of Si−H from {PhB(OxMe )2ImMes}RhH(SiH2R)CO
2
(2), is also first-order in [2] and depends on [CO]. The
equilibrium concentrations, determined over a 30 °C temperature
range, provide ΔH° = −5.5 0.2 kcal/mol and ΔS° = −16
1
cal·mol−1K−1 (for 1 ⇄ 2). The rate laws and activation parameters for oxidative addition (ΔH⧧ = 11 1 kcal·mol−1 and ΔS⧧ =
−26 3 cal·mol−1·K−1) and reductive elimination (ΔH⧧ = 17 1 kcal·mol−1 and ΔS⧧ = −10 3 cal·mol−1K−1), particularly
the negative activation entropy for both forward and reverse reactions, suggest the transition state of the rate-determining step
contains {PhB(OxMe )2ImMes}Rh(CO)2 and RSiH3. Comparison of a series of primary silanes reveals that oxidative addition of
2
arylsilanes is ca. 5× faster than alkylsilanes, whereas reductive elimination of Rh−Si/Rh−H from alkylsilyl and arylsilyl
rhodium(III) occurs with similar rate constants. Thus, the equilibrium constant Ke for oxidative addition of arylsilanes is >1,
whereas reductive elimination is favored for alkylsilanes.
insertion of metal acyl species, and low-valent, electron-rich
metal sites that react with polar substrates by oxidative
addition. Likely, low-valent metal carbonyl compounds are
present, considering that low-valent metal species readily bind
CO, formed from the decarbonylation process. Ancillary ligand
effects in these systems are, as expected, significant, but trends
are likely obscured by changing reaction conditions (reagents,
temperatures, reactants) and complex, multistep mechanisms.
In contrast, the active species are well-defined in photo-
chemically activated oxidative additions of nonpolar bonds.
Photolytic dissociation of CO is a prerequisite for reactions of
Cp*Rh(CO)2 or Tp*Rh(CO)2 (Cp* = C5Me5; Tp* = tris(3,5-
dimethyl-1-pyrazolyl)borate) with C−H or Si−H bonds.33−40
For example, Tp*Rh(CO)2 and Et3SiH react under photo-
chemical conditions to give Tp*RhH(SiEt3)CO and CO
(Figure 1).41 A sequence of photolytic CO dissociation
followed by C−H cleavage is invoked here as well as in
rhodium- or iridium-catalyzed carbonylations of benzene,
which occur under continuous irradiation.42,43 Detailed time-
resolved spectroscopic and computational studies of the
mechanism of C−H and Si−H bond activation by CpRh-
INTRODUCTION
■
Oxidative addition of nonpolar bonds (H−H, C−H, Si−H) to
low-valent rhodium and iridium metal complexes has been
widely studied because this elementary step, and the resulting
metal hydride, alkyl, or silyl species, are invoked in many
catalytic reactions including hydroformylation,1,2 hydrosilyla-
tion,3−7 and silane dehydrocoupling.8−14 Such transformations
often involve carbon monoxide directly or metal carbonyls as
catalysts, but catalytic steps involving CO dissociation are
notably absent in their proposed mechanistic cycles. Instead,
open coordination sites for oxidative addition are generated via
product-forming reductive elimination steps, and CO itself is
frequently an inhibitor in many catalytic reactions.
The strong coordination of CO must also be overcome to
affect catalytic decarbonylation chemistry. Stoichiometric
reactions of aldehydes and esters typically provide metal-
carbonyl compounds (e.g., Ni−CO or Rh−CO),15−18 and in
fact, oxidative addition of polar C−O bonds of esters to Ni(0)
is reversible upon addition of CO.18 Catalytic processes
involving decarbonylation of aldehydes,19,20 ketones,21,22 acyl
halides,17,23−26 esters,27−31 and amides32 require high temper-
atures or photochemical activation. Key intermediates
proposed in catalytic cycles for these decarbonylative trans-
formations include oxidized metal carbonyls, resulting from de-
Received: December 8, 2018
© XXXX American Chemical Society
A
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Figure 1. Photochemical and thermal pathways for oxidative addition of C−H and Si−H bonds to rhodium(I).
Scheme 1. Thermal Reactions of {PhB(OxMe )2ImMes}Rh(CO)2(1) and Primary Silanes
2
(CO),44−47 Cp*Rh(CO),37,38,48 TpRh(CO),49 Tp*Rh-
(CO),46,50−53 and Tp*Rh(CNR),40 formed via photolytic
ligand dissociation, reveal details and lifetimes of the activation
process. Interestingly, bidentate and tridentate coordination
modes of pyrazolylborate ligands modulate the electronic
structure of rhodium to affect C−H bond oxidative cleavage.
Although these studies elucidate the influence of ancillary
ligands on C−H bond oxidative addition steps, new strategies
are needed to affect both the bond activation and CO
displacement for development of new, thermal conversions.
Oxidized or electron-poor metal centers that are poorly π-
back-donating disfavor CO binding; however, such species are
also less likely to react by oxidative addition. Alternatively, a
nucleophilic metal complex could interact with the E−H bond,
and σ-coordination might reduce the π-back-donating ability to
CO and promote its dissociation. This idea implies, counter-
intuitively, that a strongly σ-donating ligand could accelerate
the dissociation of CO from an electron-rich, nucleophilic
rhodium species during a thermal oxidative addition reaction.
The trend in electron-donating ability for tridentate, fac-
coordinating ligands in rhodium dicarbonyl compounds, based
PhB(OxMe2)2ImMes (Figure 1, OxMe = 4,4-dimethyl-2-
2
oxazolinyl, ImMes = 1-(2,4,6-trimethylphenyl)imidazole).56 In
fact, PhSiH3 and {PhB(OxMe )2ImMes}Rh(CO)2 (1) react in
2
the dark to give {PhB(OxMe )2ImMes}RhH(SiH2Ph)CO (2a)
2
and CO.57
Tp*Rh(CO)2, ToMRh(CO)2, and 1 all contain mono-
anionic, multidentate ligands, which coordinate either in
fluxional bidentate or tridentate modes. The carbonyls are
photochemically labile, and all undergo C−H bond oxidative
addition upon photolysis. Both four- and five-coordinate
ground-state structures contain photolabile carbonyl ligands
and access the low coordinate states needed for C−H bond
activation.
Here we study the kinetic and thermodynamic features of
thermal Si−H bond oxidative addition to 1 and Rh−Si/Rh−H
bond reductive elimination from 2a. These investigations of
the thermal processes provide contrast with photochemical-
promoted Si−H and C−H bond oxidative additions.33−36
RESULTS AND DISCUSSION
■
Synthesis of Rhodium Silyl Compounds. The reaction
of 1 and 3 equiv of PhSiH3 gives 2a over 18 h in benzene at
room temperature in 94% isolated yield.57 Similarly, 1 and
primary alkylsilanes (R = C6H13, C12H25) or arylsilanes (R = p-
on CO stretching frequencies, is Tp < Tp* < ToM (ToM
=
tris(4,4-dimethyl-2-oxazolinyl)phenylborate, Figure 1);54,55
however, ToMRh(CO)2 does not appear to react with
hydrocarbons or silanes under thermal conditions, and we
concluded that the tris(oxazolinyl)borate ligand was not
sufficiently electron donating for the rhodium to undergo the
desired oxidative addition. In order to further increase the
electron-donating ability of the fac-coordinating ligand system,
we replaced one oxazoline in ToM with a N-heterocyclic
carbene to give the bis(oxazolinyl)(NHC)borate ligand
MeC6H4, p-MeOC6H4) react to give {PhB(OxMe )2ImMes}-
2
RhH(SiH2R)CO (R = C6H13 (2b), C12H25 (2c), p-MeC6H4
(2d), p-MeOC6H4 (2e); Scheme 1). The reactions of
alkylsilanes and 1 are slower than the corresponding reactions
of arylsilanes under equivalent reaction conditions. For
example, the reaction of 1 and hexylsilane (10 equiv) in
benzene requires 1 d at room temperature to form 2b
B
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quantitatively, whereas the reaction of PhSiH3 under
equivalent conditions affords 2a quantitatively after 6 h.
In the solid-state structure of 2d (Figure 2), the Rh1−N1
interatomic distance trans to the hydride is shorter by 0.05 Å
undergo rearrangement to Ph2SiH2 or R2SiH2 at 60 °C, instead
giving the desired rhodium silyl hydride product selectively.
Kinetic Studies of Oxidative Addition. Initial kinetic
studies, in which concentrations of PhSiH3 and 1 were
1
monitored by H NMR spectroscopy, suggested first-order
dependence of reaction rate on [PhSiH3] and [1]. Unfortu-
nately, reactions in sealed NMR tubes do not result in
complete conversion, even with excess PhSiH3. For example,
15 equiv of PhSiH3 (140 mM) vs 1 (9 mM) gives 92% yield of
2a after 70 min (1:2a = 0.085:1). Although the concentration
of 1 follows an exponential decay over the first few half-lives,
the composition of the reaction mixture changes more slowly
in the later stages of the reaction than expected for a pseudo-
first-order reaction that should proceed to completion.
Moreover, a plot of pseudo-first-order rate constants vs
[PhSiH3] provides only a rough trend, and considerable
scatter suggests additional factors influence the rate.
A unimolecular decomposition via reductive elimination of
the rhodium silyl hydride is ruled out because compound 2a is
isolable and persistent in pure form. In contrast, the reaction of
2a and CO (1 atm) results in ca. 30% conversion to 1 and
PhSiH3 at room temperature after 1 d. This observation
indicates that oxidative addition of Si−H to 1 is reversible in
the presence of CO. Thus, reproducible kinetic measurements
on reactions of 1 might be complicated by slow and variable
loss of CO into the NMR tube headspace. An additional
complication to studying the CO-promoted reductive elimi-
Figure 2. Thermal ellipsoid plot of {PhB(OxMe )2ImMes}RhH-
2
nation from 2a involves the formation of PhB(OxMe )2ImMesH
2
(SiH2C6H4Me)CO (2d) at 35% probability. One of two positions
of the disordered tolyl group is shown. H atoms bonded to Rh1 and
Si1 were located objectively in the difference Fourier map, refined
isotropically, and were included in the representation. Selected
interatomic distances (Å): Rh1−Si1, 2.336(1); Rh1−C11, 2.070(4);
Rh1−C36, 1.888(3); Rh1−H1r, 1.58(4); Rh1−N1, 2.197(3); Rh1−
N2, 2.242(3).
as a side product under conditions with higher CO pressures
or high temperature (>60 °C). In contrast, the reaction
produces only 1 and PhSiH3 under ambient conditions.
Apparently, the number of CO ligands coordinated to Rh
increases at higher pressure, leading to reductive elimination of
C−H rather than Si−H. Note that the harsher conditions
required for C−H reductive elimination indicate that the Si−H
reductive elimination is kinetically preferred.
Solution-phase IR and UV−vis spectroscopies were tested as
alternative methods to study the reactions, with the latter
method appearing to be most promising. UV−vis spectroscopy
provides the advantage that the disappearance of 1 could be
monitored in sealed cuvettes with negligible headspaces,
eliminating one of the issues associated with the NMR kinetic
studies. In a first experiment, spectra (250−700 nm) of a
reaction mixture of 1 and PhSiH3 were acquired at 15 min
intervals over the reaction (Figure 3). The signal at 385.5 nm
systematically decreased, while the center of the band at 301
nm blue-shifted and appeared non-Gaussian in the final
spectrum of 2a. Subsequent kinetic experiments monitored the
absorption at 385.5 nm. While photolysis of 1 in benzene
than the Rh1−N2 distance of the oxazoline trans to the silyl
group. The CO ligand and the NHC group are trans. This
configuration and the shorter Rh−N interatomic distance trans
to H are consistent across the solid-state structures of 2a57 and
2b (see Figure S17). In contrast, the reaction of {PhB-
(OxMe )2ImMes}Ir(CO)2 and benzene, which occurs under
2
photochemical conditions, results in CO trans to an oxazoline
and the phenyl trans to the NHC.57 These distinct
configurations of the two products suggest distinct mechanisms
for thermal and photochemical processes.
Contamination by residual arylsilane hinders reproducible
isolation of 2d and 2e as pure species. These compounds are
instead spectroscopically characterized in the presence of
RSiH3. Attempts to increase the rate of oxidative addition of p-
MeC6H4SiH3 or p-MeOC6H4SiH3, by performing the reactions
at 60 °C, result in organosilane redistribution to give (p-
MeC6H4)2SiH2 or (p-MeOC6H4)2SiH2. Neither triarylsilanes
nor the SiH4 product expected from redistribution were
detected by NMR spectroscopy or GC-MS; however, addi-
tional SiH and RhH signals suggest that these species, if
formed, might have undergone further reaction. The majority
of redistribution products appear in the latter stages of the
reaction after almost all of 1 is consumed, suggesting that 2d
and 2e are involved in redistribution. This process is faster for
p-MeOC6H4SiH3 than for p-MeC6H4SiH3, and only the
oxidative addition pathway is observed at room temperature.
Interestingly, neither PhSiH3 nor the alkylsilanes appear to
affords the cyclometalated {κ4-PhB(OxMe )2ImMes′CH2}RhH-
2
(CO) resulting from photolytic CO dissociation followed by
CH bond oxidative addition of a mesityl methyl,57 that process
was not observed (from absorption of the UV light) in the
spectrophotometer. That is, a series of UV−vis spectra for 1
acquired under otherwise equivalent conditions, but without
PhSiH3, were all identical to the first spectrum.
Under conditions employing large excesses of PhSiH3
(ranging from 25 to 140 equiv vs 1) for reactions monitored
at 296 K, plots of [1] vs time followed an exponential decay for
3 half-lives, indicating first-order dependence on [1]. Under
the conditions of these experiments, particularly with dilute [1]
and high [PhSiH3], the final concentration of 1 approaches
zero (Figure 4). Nonlinear least-squares fits were superior for
C
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app
Table 1. Second-Order Rate Constants k1 from Pseudo-
a
First-Order Kinetic Studies
temp. (K) kH M−1 s−1 (×10−2
)
kD M−1 s−1 (×10−2
app
)
kH/kD
app
b
296.0
297.4
308.3
312.3
316.8
1.1 0.1
1.2
2.0 0.1
2.9 0.3
3.7 0.1
4.2 0.4
1.1
1
0.9
0.9 0.1
1.0 0.1
0.99 0.02
0.9 0.1
b
1.3 0.1
2.3 0.1
3.0 0.1
3.7 0.3
5.0 0.1
322.3
a
app
k1 is the rate constant for the forward, oxidative addition of
b
PhSiH3 to 1, determined at [CO]ini = 0. Calculated from
corresponding Eyring plots.
kcal·mol−1 and ΔS⧧H‑app = −30 3 cal·mol−1K−1 (Figure S24).
For comparison, the ΔS⧧ for oxidative addition of Si−H bonds
in tertiary silanes to photolytically generated CpMn(CO)2
range from −6 to −10 cal·mol−1K−1;58 however, these values
vary considerably with the cyclopentadiene-based ancillary
ligand.59
The same approach provides the second-order rate constants
for the addition of 1 and PhSiD3. The activation parameters,
determined from an Eyring plot, are ΔH⧧D‑app = 10 1 kcal·
mol−1 and ΔS⧧D‑app = −34 4 cal·mol−1K−1. These two sets of
values are experimentally indistinguishable, thus kH/kD ∼ 1 for
all accessible reaction temperatures.
Figure 3. Electronic absorption spectra of 1 ([1]ini = 2.7 × 10−4 M)
during its reaction with PhSiH3 (1.1 × 10−3 M) at room temperature
in benzene acquired at 15 min intervals. PhSiH3 does not absorb at
>250 nm.
Equilibrium Kinetic Studies. Additional evidence that
oxidative addition of phenylsilane is reversible in the presence
of CO comes from the reaction with PhSiH3 in a CO-saturated
benzene solution (7.3 mM).60 Under otherwise equivalent
conditions, the reaction shows an apparent decrease in rate and
a nonzero final concentration of 1 (Figure 5).
Figure 4. Plots of [1] vs time in its reaction with PhSiH3 in benzene
at 296 K, monitored at 385.5 nm, with [PhSiH3] = (a) 4.0, (b) 5.8,
(c) 10.0, and (d) 25.5 mM. Nonlinear least-squares regression fits to
exponential decay curves [Rh]t = [Rh]ini e−kt confirm first-order
dependence.
experiments with higher concentrations of PhSiH3, while some
deviation was observed at longer reaction times with lower
[PhSiH3]ini. This observation suggests that at longer reaction
times, [CO] increases, and the reverse reaction contributes to
the observed rate constant (see below for studies with nonzero
initial CO concentrations). Thus, the initial portion of the time
course approximates the reaction as irreversible under
conditions where [CO]ini = 0 M.
Figure 5. Plots of [1] vs time for its reaction with PhSiH3 (260 equiv)
in benzene saturated with (a) ca. 35 equiv of CO and (b) without CO
at 296 K, measured by UV−vis spectroscopy at 385.5 nm. [CO]ini
7.3 mM, [1]ini = 0.2 mM, [PhSiH3]ini = 52.0 mM.
=
Plots of kobs vs [PhSiH3] provide a forward rate constant kapp
(given in Figure S21, constrained by the approximation that
the rate of reductive elimination is negligible at [CO]ini = 0
M). In addition, a plot of log[PhSiH3] vs log(kobs) is linear
with a slope of 1 (Figure S23), indicating first-order
dependence on phenylsilane concentration. The rate constants
were measured over five temperatures (up to 322 K, Table 1)
Clearly, mechanisms of both forward and reverse reactions
are of interest. Note that for a first-order reaction that
approaches equilibrium, the observed rate constant (i.e., the
coefficient of the exponential describing time-dependence of
the reactant’s concentration) is the sum of forward and reverse
rate constants (ke = k1 + k−1).61 In the present case, the
observed pseudo-first-order rate constant (ke) is equal to the
sum of the oxidative addition and reduction elimination rate
to determine the activation parameters of ΔH⧧H‑app = 11
1
D
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constants, multiplied by concentrations of flooded reagents,
respectively (eq 1).
l
o
|
o
1
1
[1]e
o
o
ke = k1[PhSiH3]e[1]em
+
}
o
o
o
o
[PhSiH3]e
n
~
l
|
o
o
1
1
o
o
+ k−1[CO]e × [2a]em
+
}
o
o
o
o
[CO]e
[2a]e
(1)
n
~
Under high [PhSiH3] and [CO] (flooding conditions), eq 1
simplifies into eq 2, where the observed rate constant ke
depends on both [PhSiH3]e and [CO]e (which are defined
by their initial concentrations). The forward and reverse rate
constants are related to the equilibrium constant Ke = k1/k−1.
k
ke = k1[PhSiH3]e + k−1[CO]e = k1[PhSiH3]e + 1 [CO]e
Ke
(2)
To determine Ke, 1 (0.1−0.2 mM) and PhSiH3 (4.8−8 mM)
are allowed to react in CO-saturated benzene solution (7.3
mM) over 1−2 days, until the absorbance of 1 becomes static
and corresponds to [1]e. [2a]e is determined from the
Figure 6. Eyring plots (ln(k/T) vs 1/T) for (a) k1 representing
oxidative addition and (b) k−1 representing reductive elimination.
1
relationship [1]ini − [1]e. At 297 K, Ke is 2.8
0.2 (from
least on the NMR time scale), and its H NMR spectrum
four independent measurements). The equilibrium constants
were measured from 297 to 328 K, and a van’t Hoff plot
provides ΔH° = −5.5 0.2 kcal/mol and ΔS° = −16 1 cal·
mol−1K−1 (Figure S25). The van’t Hoff plot also allows the
prediction of equilibrium constants for other temperatures.
The concentration of 1 vs time was measured under
conditions of excess PhSiH3 and excess CO (7.3 mM).
Nonlinear least-squares regression of the data to eq 3 provides
ke.
contains distinct resonances for the oxazoline trans to hydride
and silyl, as noted above, suggesting that five-coordinate
fluxional intermediates, possibly from reductive coupling of
Si−H or dissociation of an oxazoline, are not involved. From
these data, a probable rate law for reductive elimination is
second-order (∝ [2a][CO]). These results, and the principle
of microscopic reversibility, indicate that the composition of
the transition state for forward and reverse processes is
[{PhB(OxMe )2ImMes}Rh(CO)2·PhSiH3]. Thus, the reverse
2
pathway involves coordination of CO to 2a as a first step.
Rates and Equilibrium Constants for Organosilane
Addition. The equilibrium constant Ke, forward rate constant
k1, and reverse rate constant k−1 are determined for the
reaction of 1 with four primary silanes (Table 2). Comparison
of Ke values reveals that oxidative addition of arylsilanes is
more favorable than the corresponding reaction of alkylsilanes.
The forward rate constant (k1) is nearly five times faster for
arylsilanes than for alkylsilanes, whereas the rate constants for
the reverse reaction are similar across all silanes. Because the
reverse rate constants are relatively similar, the difference in
equilibrium constant between alkylsilanes and arylsilanes
results from the difference in forward rate constant. Note
that only two of the three variables (Ke, k1, and k−1) are
independent.
The rate constant for oxidative addition of alkylsilanes is
smaller than that for reductive elimination, whereas reductive
elimination of arylsilanes occurs with a lower rate constant
than the oxidative addition. Even though this thermodynamic
assessment indicates that compounds 2d and 2e are favored
while 2b and 2c are unfavored with respect to 1 (without
excess reagent to perturb the equilibrium), somewhat ironi-
cally, the latter compounds are the isolable species. Clearly, the
unfavorable thermodynamic bias may be overcome under
synthetic conditions by removing CO from the system.
[1] = [1]e + ([1]0 − [1]e)e−k t
e
(3)
The rate constants k1 and k−1 are determined using eq 2 and ke,
Ke, [PhSiH3]e and [CO]e values. Measurements of [1] vs time
from 297 to 328 K (four measurements at each temperature)
provide temperature-dependent ke (Table S1). These rate
constants, combined with the Ke from the van’t Hoff analysis,
provide k1 and k−1 over a 30 K range. Here, we note that the
rate constant k1
of 0.01 M−1 s−1 (calculated from the
296 K
Eyring equation and van’t Hoff plot for a reaction occurring at
296 K) and the kapp (0.012 M−1 s−1, measured experimentally
at 296 K) are in good agreement, indicating that the
approximate rate constant at [CO]ini = 0 is a reasonable
estimate of k1.
The activation parameters for oxidative addition of 1 and
PhSiH3 are determined to be ΔH⧧= 11 1 kcal·mol−1 and
ΔS⧧ = −26 3 cal·mol−1·K−1 (Figure 6). These values are also
equivalent to those determined above from rate constants kapp
in which the reverse direction is neglected at [CO]ini = 0. The
second-order rate law indicates that 1 and PhSiH3 are present
in that transition state, while this negative entropy of activation
suggests (but does not prove) that CO (which ultimately
dissociates) is also present in the transition state as the rate-
controlling step occurs. The reverse reaction also features a
negative entropy of activation (ΔH⧧ = 17 1 kcal·mol−1 and
ΔS⧧ = −10
3 cal·mol−1K−1). In addition, the rhodium
CONCLUSION
tolylsilyl (2d, see below) is not detected in NMR spectra of a
mixture of 2a and p-tolylsilane (as long as the solution is CO-
free). Thus, the reductive Si−H coupling and silane formation
from 2 requires a molecule of CO to interact with the rhodium
center. Moreover, compound 2a is stereochemically rigid (at
■
Comparison of the forward and reverse rate constants between
reaction of arylsilanes and alkylsilanes as well as the rate laws
and activation parameters for PhSiH3 oxidative addition and
reductive elimination provide considerable insight into the
E
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Table 2. Ke, k1, and k−1 for Reaction of 1 and RSiH3 Measured at 297 °C
RSiH3/2
Ke
k1 (M−1 s−1) (×10−2
)
k−1 (M−1 s−1) (×10−2
)
n-hexylsilane/2b
0.53 0.01
0.48 0.01
2.8 0.3
0.23 0.01
0.25 0.02
1.25
0.43 0.01
0.53 0.05
0.49
n-dodecylsilane/2c
a
phenylsilane /2a
p-methoxyphenylsilane/2e
p-tolylsilane/2d
2.1 0.1
4.5 0.1
1.34 0.02
1.4 0.1
0.64 0.01
0.32 0.02
a
Values for PhSiH3 were determined from Eyring and van’t Hoff plots.
pathway and mechanisms. First, reductive elimination (the
back reaction) likely follows a second-order rate law. The
dissociation of silane from rhodium does not occur in the
absence of free CO, the rhodium center in 2 is stereochemi-
cally rigid, and silyl/silane exchange is not observed. Thus, two
CO groups are present in the transition state of the rate-
limiting step on the pathway for reductive elimination (eq 6 of
Scheme 2, in reverse). Alternative mechanisms, in which one
oxazoline by the incoming CO. This step is the first step in the
overall pathway from 2 back to 1.
This interpretation limits the possible mechanisms of the
oxidative addition (forward) pathway (Scheme 2). In oxidative
addition of Si−H to 1, two CO ligands are present up to the
transition state of the rate-limiting step, at which point CO
dissociates from rhodium, as required by the principle of
microscopic reversibility. Thus, low coordinate species,
analogous to photochemically generated κ2-Tp*RhL or κ3-
Tp*RhL (L = CO, CNR), are never accessed in this
transformation. In addition, the Si−H bond is cleaved during
or prior to the rate-limiting step in the reaction of 1 and
silanes. The evidence for this aspect of the mechanism includes
the second-order rate law which contains first-order depend-
ence on silane concentration. The activation entropy (ΔS⧧ =
−26 3 cal·mol−1·K−1) is more negative than for the reverse
direction, which is consistent with the loss of entropy resulting
from creation of the rhodium-silane interaction. Finally, the
rate constants for oxidative addition to 1 are 5× greater with
arylsilanes than alkylsilanes, indicating that the relative energy
of the starting material vs that of the rate-determining
transition state is affected by the nature of the silane. Thus,
we attribute the change in rate of oxidative addition for 1 and
silanes to large changes in the energy of the silane (variation in
the ground state). This interpretation is consistent with the
similar rates for reductive elimination over the series 2a−e,
which indicates the energy differences between 2 and the rate-
limiting transition for reductive elimination are similar for all
the silyl species.
Scheme 2. Proposed Pathway for Oxidative Addition of Si−
H Bonds under Thermal Conditions
In contrast, we note that kH/kD is ∼ 1. The most
straightforward interpretation of these data is that the Si−H
is broken prior to the rate-determining step. A second
possibility is that Si−H bond cleavage occurs during the
rate-determining step, and an inverse equilibrium isotope effect
from the formation of a silane-rhodium σ-complex balances a
normal kinetic isotope effect for Si−H oxidative cleavage.62
The data cannot unambiguously distinguish these alternatives,
but we favor the former pathway of Scheme 2 because the
observed kinetic isotope effect is unity over a range of
temperatures, such that coincidental canceling of equilibrium
and kinetic isotope effects is less likely, and this scheme
provides the simplest kinetically consistent mechanism.
The thermodynamic and kinetic data are entirely incon-
sistent with a mechanism resembling the photochemical
activation and E−H bond oxidative addition chemistry of the
classical cyclopentadienyl and tris(pyrazolyl)borato rhodium
and iridium species. The newly proposed mechanism invokes
associative steps to access the trivalent rhodium intermediate,
which is then susceptible to displacement of CO by a chelating
oxazoline. This work shows that thermal pathways for catalytic
functionalization are plausible, even in the presence of CO, but
such reactions still face considerable challenges due to kinetic
and thermodynamic constraints on bond activation. We also
oxazoline dissociates from 2 to give a high-energy 16-electron
rhodium(III) species that is trapped by CO or undergoes Si−
H bond reductive elimination, are ruled out by the dependence
on CO.
Two additional observations suggest that coordination of
CO to 2 is the rate-limiting step of the reductive elimination
process. In this reaction, ΔS⧧ is negative (−10
3 cal·
mol−1K−1), which is consistent with increased order in the
transition state. Loss of the H-SiH2R group during or prior to
the rate-limiting step would result in a positive entropy term.
Note that ΔS for the overall reductive elimination of PhSiH3
from 2a is positive, suggesting that PhSiH3 dissociation from
2a occurs after the rate-controlling step. In addition, the rate
constants are similar for the reductive elimination reactions of
2a−e, where the alkylsilyl and arylsilyl ligand is varied. A rate-
determining step involving either Rh−Si bond cleavage, and
especially Si−H bond formation, would be expected to vary
significantly with the silyl substituent. On the basis of these
ideas, we conclude that the rate-limiting step for the reductive
elimination pathway involves substitution of a coordinated
F
DOI: 10.1021/acs.inorgchem.8b03425
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
in benzene to give a pale yellow solution. The resulting solution was
allowed to stir for 24 h. The volatile materials were evaporated in
vacuo resulting in a white solid which was washed with pentane to
find that CO dissociation from a low-valent metal complex is
not a necessary prerequisite for oxidative addition of nonpolar
substrates. The presence of coordinating groups, either as part
of polar substrates, ancillary ligands, or as σ-complexes, may
also facilitate coordination, oxidative addition, and CO
dissociation during catalytic conversions.
give pure {PhB(OxMe )2ImMes}RhH(SiH2C12H25)CO(65 mg, 0.08
2
mmol, 50%). 1H NMR(benzene-d6, 600 MHz): δ 8.41 (d, 3JHH = 7.8
3
Hz, 2 H, o-BC6H5), 7.54 (t, JHH = 7.2 Hz, 2 H, m-BC6H5), 7.4 (t,
3JHH = 7.2 Hz, 1 H, p-BC6H5), 6.84 (s, 1 H, m-C6H2Me3), 6.80 (s, 1
H, m-C6H2Me3), 6.54 (d, 3JHH = 1.6 Hz, 1 H, N2C3H2Mes), 5.94 (d,
3JHH = 1.8 Hz, 1 H, N2C3H2Mes), 4.26 (t, 3JHH = 7.2 Hz, 1JSiH = 160
EXPERIMENTAL SECTION
■
3
1
General. All reactions were performed under a dry argon
atmosphere using standard Schlenk techniques or under a nitrogen
atmosphere in a glovebox. Benzene, toluene, methylene chloride,
pentane, and tetrahydrofuran were dried and deoxygenated using an
IT PureSolv system. Benzene-d6 was heated to reflux over Na/K alloy
Hz, 1 H, SiH), 3.81 (t, JHH = 6.6 Hz, JSiH = 180 Hz, 1 H, SiH),
2
3.64−3.57(m, 3 H, CNCMe2CH2O), 3.35 (d, JHH = 8.4 Hz, 1 H,
CNCMe2CH2O), 2.26 (s, 3 H, p-C6H2Me3), 2.09 (s, 3 H, o-
C6H2Me3), 2.03 (s, 3 H, o-C6H2Me3), 1.58 (br, 2 H, SiCH2CH2),
1.33−1.29 (br, 21 H, CH2CH2CH2), 1.23 (s, 3 H, CNCMe2CH2O
trans to H), 1.16 (s, 3 H, CNCMe2CH2O trans to H), 1.10 (s, 3 H,
CNCMe2CH2O trans to Si), 0.99 (s, 3 H, CNCMe2CH2O trans to
Si), 0.93 (t, 3JHH = 7.2 Hz, 3 H, CH2CH3), 0.44 (m, 1 H, SiH2CH2),
and vacuum transferred. [Rh(μ-Cl)(CO)2]2,63 {PhB(OxMe )2ImMes}-
2
57
Rh(CO)2 (1), {PhB(OxMe )2ImMes}RhH(SiH2Ph)CO, and p-
2
methoxyphenylsilane were synthesized according to the literature
procedures.64 Phenylsilane, n-hexylsilane, n-dodecylsilane, and p-
tolylsilane were synthesized by reducing corresponding trichlorosi-
lanes with LiAlH4. PhSiD3 was synthesized from PhSiCl3 and LiAlD4.
Potassium benzyl was synthesized by reacting potassium tert-butoxide
with n-BuLi in toluene.65
1
3
−0.07 (m, 1 H, SiH2CH2), −13.68 (dd, JRhH = 22.2 Hz, JHH = 1.8
Hz, 1 H, RhH). 13C{1H} NMR (benzene-d6, 150 MHz): δ 195.06 (d,
1JRhC = 52.8 Hz, 2C−N2C3H2Mes), 185.79 (br, CNCMe2CH2O),
179.23 (d, 1JRhC = 40.7 Hz, CO), 138.49 (p-C6H2Me3), 137.60 (ipso-
C6H2Me3), 136.70 (o-C6H2Me3), 136.66 (o-C6H2Me3), 136.39 (o-
BC6H5), 129.59 (m-C6H2Me3), 129.52 (m-C6H2Me3) 127.51 (m-
BC6H5), 127.04 (p-BC6H5), 124.61 (4,5-CN2C3H2Mes), 121.12 (4,5-
CN2C3H2Mes), 80.40 (CNCMe2CH2O), 80.39 (CNCMe2CH2O),
68.79 (CNCMe2CH2O), 66.64 (CNCMe2CH2O), 34.00 (CH2),
32.38 (CH2), 30.37 (CH2), 30.31 (CH2), 30.26 (CH2), 30.19
(CH2), 29.94 (CH2), 29.87 (CH2), 28.37 (CNCMe2CH2O), 28.08
(CNCMe2CH2O), 28.02 (CNCMe2CH2O), 27.03 (CNCMe2CH2O),
23.16 (CH2), 21.18 (p-C6H2Me3), 19.53 (o-C6H2Me3), 19.38 (o-
C6H2Me3), 18.02 (CH2), 18.01 (CH2), 14.41 (CH3). 15N{1H} NMR
(benzene-d6, 61 MHz): δ −160 (CNCMe2CH2O trans to Si), −172
(CNCMe2CH2O trans to H), −175 (N2C3H12Mes), −188
(N2C3H2Mes). 11B NMR (benzene-d6, 192 MHz): δ −9.7. 29Si{1H}
NMR (benzene-d6, 119 MHz): δ −18.9. IR (KBr, cm−1): 2957, 2922,
2852, 2149, 2044, 2013, 1595, 1463, 1365, 1316, 1183, 1160, 968.
1H, 13C{1H}, 11B, and 15N spectra were collected on Bruker Avance
III 600 or AVNEO 400 MHz NMR spectrometers. NMR signals (1H,
13C, and 15N) were assigned based on COSY, HMQC, and HMBC
experiments. 15N chemical shifts were determined by 1H−15N HMBC
experiments. 15N chemical shifts were initially referenced to NH3 and
recalibrated to the CH3NO2 chemical shift scale by adding −381.9
ppm. Infrared spectra were recorded on a Bruker Vertex spectrometer.
Elemental analyses were performed using a PerkinElmer 2400 Series
II CHN/S in the Iowa State Chemical Instrumentation Facility.
{PhB(OxMe )2ImMes}RhH(SiH2C6H13)CO (2b). n-Hexylsilane (185
2
mg, 1.6 mmol) was added to a solution of 1 (100 mg, 0.158
mmol) in benzene to give a pale yellow solution. The resulting
solution was allowed to stir for 24 h. The volatile materials were
evaporated in vacuo giving {PhB(OxMe )2ImMes}RhH(SiH2C6H13)CO
2
(88 mg, 0.123 mmol, 78%). 1H NMR (benzene-d6, 600 MHz): δ 8.40
{PhB(OxMe )2ImMes}RhH(SiH2C6H4Me)CO (2d). p-Tolylsilane (195
2
3
3
(d, JHH = 7.8 Hz, 2 H, o-BC6H5), 7.54 (t, JHH = 7.8 Hz, 2 H, m-
mg, 1.6 mmol) was added to a solution of 1 (100 mg, 0.158 mmol) in
benzene to give a pale yellow solution. The resulting solution was
allowed to stir for 24 h. The volatile materials were evaporated in
3
BC6H5), 7.39 (t, JHH = 7.3 Hz, 1 H, p-BC6H5), 6.83 (s, 1 H, m-
3
C6H2Me3), 6.78 (s, 1 H, m-C6H2Me3), 6.53 (d, JHH = 1.6 Hz, 1 H,
N2C3H2Mes), 5.94 (d, 3JHH = 1.6 Hz, 1 H, N2C3H2Mes), 4.91 (t, 3JHH
vacuo giving {PhB(OxMe )2ImMes}RhH(SiH2C6H4Me)CO with some
2
= 6.8 Hz, 1JSiH = 168 Hz, 1 H, SiH), 4.43 (t, 3JHH = 6.8 Hz, 1JSiH = 180
residual p-tolylsilane. This mixture was characterized by NMR and IR
spectroscopy; the residual silane hindered 13C{1H} NMR assignments
in the aryl region. An X-ray quality crystal was obtained from a
pentane solution; however, this approach was not reliable for
2
Hz, 1 H, SiH), 3.60 (m, 3 H, CNCMe2CH2O), 3.36 (d, JHH = 8.3
Hz, 1 H, CNCMe2CH2O), 2.23 (s, 3 H, p-C6H2Me3), 2.09 (s, 3 H, o-
C6H2Me3), 2.02 (s, 3 H, o-C6H2Me3), 1.55 (br, 2 H, SiCH2CH2),
1.33−1.40 (br, 6 H, CH2CH2CH2), 1.23 (s, 3 H, CNCMe2CH2O
trans to H), 1.16 (s, 3 H, CNCMe2CH2O trans to H), 1.10 (s, 3 H,
CNCMe2CH2O trans to Si), 0.99 (s, 3 H, CNCMe2CH2O trans to Si),
0.95 (t, 3JHH = 7 Hz, 3 H, CH2CH3), 0.41 (m, 1 H, SiH2CH2), −0.12
1
purification from residual p-tolylsilane. H NMR (benzene-d6, 600
3
3
MHz): δ 8.43 (d, JHH = 7.8 Hz, 2 H, o-BC6H5), 7.56 (t, JHH = 7.2
Hz, 2 H, m-BC6H5), 7.52 (d, 3JHH = 7.8 Hz, 2 H, o-C6H4Me), 7.41 (t,
3JHH = 7.2 Hz, 1 H, p-BC6H5), 6.95 (d, JHH = 7.8 Hz, 2 H, m-
3
(m, 1 H, SiH2CH2), −13.70 (dd, 1JRhH = 22.4 Hz, 3JHH = 1.5 Hz, 1 H,
C6H4Me), 6.55 (d, 3JHH = 1.8 Hz, 1 H, N2C3H2Mes), 6.54 (s, 1 H, m-
1
RhH). 13C{1H} NMR (benzene-d6, 150 MHz): δ 195.05 (d, JRhC
=
3
1
C6H2Me3), 6.41 (s, 1 H, m-C6H2Me3), 5.94 (d, JHH = 1.8 Hz, 1 H,
51.9 Hz, 2C-N2C3H2Mes), 179.24 (d, JRhC = 40.5 Hz, CO), 138.49
(br, ipso-BC6H5), 137.58 (p-C6H2Me3), 136.70 (o-C6H2Me3), 136.64
(o-BC6H5), 136.36 (o-C6H2Me3), 129.58 (m-C6H2Me3), 129.51 (m-
C6H2Me3), 127.50 (m-BC6H5), 127.03 (p-BC6H5), 124.60 (4,5C−
N2C3H2Mes), 121.12 (4,5C−N2C3H2Mes), 80.40 (CNCMe2CH2O),
80.39 (CNCMe2CH2O), 68.79 (CNCMe2CH2O), 66.64
(CNCMe2CH2 O), 33.64 (CH2), 32.32 (CH2), 29.88
(CNCMe2CH2O), 28.37 (CNCMe2CH2O), 28.08 (CNCMe2CH2O),
27.03 (CNCMe2CH2O), 23.29 (CH2), 21.16 (p-C6H2Me3), 19.52 (o-
C6H2Me3), 19.37 (o-C6H2Me3), 18.01 (CH2), 17.99 (CH2), 14.50
(CH3). 15N{1H} NMR (benzene-d6, 61 MHz): δ −160
(CNCMe2CH2O trans to Si), −172 (CNCMe2CH2O trans to H),
−188 (N2C3H2Mes). 11B NMR (benzene-d6, 192 MHz): δ −9.7.
29Si{1H} NMR (benzene-d6, 119 MHz): δ −19.2. IR (KBr, cm−1):
3135, 3043, 2959, 2923, 2733, 2359, 2279, 2041, 2011, 1652, 1595,
1458, 1275, 967, 820. Anal. calcd for C35H50BN4O3RhSi: C, 58.66; H,
7.03; N, 7.82. Found: C, 58.68; H, 7.21; N, 7.61
N2C3H2Mes), 4.91 (t, 2JHH = 4.8 Hz, 1JSiH = 170 Hz, 1 H, SiH), 4.45
2
1
(d, JHH = 4.8 Hz, JSiH = 186 Hz, 1 H, SiH), 3.57−3.65 (m, 3 H,
CNCMe2CH2O), 3.37 (d, 2JHH = 8.4 Hz, 1 H, CNCMe2CH2O), 2.18
(s, 3 H, p-C6H2Me3), 2.03 (s, 3 H, o-C6H2Me3), 2.00 (s, 3 H, o-
C6H2Me3), 1.90 (s, 3 H, p-C6H4Me), 1.16 (3 H, CNCMe2CH2O),
1.16 (3 H, CNCMe2CH2O), 1.08 (3 H, CNCMe2CH2O), 1.02 (3 H,
1
3
CNCMe2CH2O), −13.22 (dd, JRhH = 21.6 Hz, JHH = 2.4 Hz, 1 H,
RhH). 13C{1H} NMR(benzene-d6, 150 MHz): δ 194.48 (d, JRhC
=
1
1
52.9 Hz, 2C-N2C3H2Mes), 178.40 (d, JRhC = 40.8 Hz, CO), 138.29
(ipso-SiC6H4Me), 136.96 (p-C6H2Me3), 136.68 (ipso-C6H2Me3),
136.67 (o-BC6H5), 129.34 (o-SiC6H4Me), 129.31 (m-SiC6H4Me),
127.54 (m-C6H2Me3), 127.07 (m-C6H2Me3), 124.60 (4,5C−
N2C3H2Mes), 121.50 (4,5C−N2C3H2Mes), 80.49 (CNCMe2CH2O),
80.31 (CNCMe2CH2O), 68.80 (CNCMe2CH2O), 66.70
(CNCMe2CH2O), 28.38 (CNCMe2CH2O), 28.07 (CNCMe2CH2O),
27.78 (CNCMe2CH2O), 27.03 (CNCMe2CH2O), 21.48 (p-
C6H2Me3), 21.04 (p-SiC6H4Me), 19.49 (o-C6H2Me3), 19.08 (o-
C6H2Me3). 15N{1H} NMR (benzene-d6, 61 MHz): δ −161
{PhB(OxMe )2ImMes}RhH(SiH2C12H25)CO (2c). n-dodecylsilane (318
2
mg, 1.59 mmol) was added to a solution of 1 (100 mg, 0.159 mmol)
G
DOI: 10.1021/acs.inorgchem.8b03425
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
(CNCMe2CH2O trans to Si), −172, (CNCMe2CH2O trans to H),
−174 (N2C3H2Mes), −188 (N2C3H2Mes). 11B NMR (benzene-d6,
192 MHz): δ −9.8. IR (KBr, cm−1): 2962, 2923, 2855, 2279, 2104,
2016, 1917, 1652, 1605.
freeze−pump−thaw cycles and then saturated with CO by stirring
under a CO atmosphere (298 K, 1 atm). The solutions were
transferred to nitrogen-flushed and dried UV−vis cuvettes using air-
free syringe technique. These samples were used for the kinetic study,
following the procedure above, to determine ke by nonlinear least-
squares analysis. The series of experiments, combined with Ke
(determined below), provide forward (k1) and reverse (k−1) rate
constants. This procedure was repeated for the three other
temperatures (308.4, 318.4, 328.4 K). For the silane p-tolylsilane, p-
methoxysilane, hexylsilane, and dodecylsilane, this procedure was
applied at room temperature to determine forward and reverse rate
constants.
Equilibrium Constant Measurement. Reactions of 1 and PhSiH3
(3.2 mM, 3.8 mM, 4 mM, 11.9 mM), in CO-saturated benzene
solution, were monitored by UV−vis, as described above, until the
reaction mixture reached equilibrium. [1] was determined from Beer’s
law, [2a]e was determined from difference between [Rh]total and [1]e,
and [PhSiH3]e and [CO]e were approximated to be equal to their
initial concentration (in large excess of [Rh]total). This procedure was
repeated for other temperatures (296, 301.5, 318.4, 328.4 K) to
provide data for a van’t Hoff plot to determine ΔH and ΔS.
{PhB(OxMe )2ImMes}RhH(SiH2C6H4OMe)CO (2e). p-Methoxyphenyl-
2
silane (220 mg, 1.6 mmol) was added to a solution of 1 (100 mg,
0.158 mmol) in benzene to give a pale yellow solution. The resulting
solution was allowed to stir for 24 h. The volatile materials were
evaporated in vacuo giving {PhB(OxMe )2ImMes}RhH-
2
(SiH2C6H4OMe)CO with some residual p-methoxyphenylsilane.
This mixture was characterized by NMR and IR spectroscopy; the
residual silane and its slow catalyzed redistribution hindered 13C{1H}
NMR assignments in the aryl region. 1H NMR (benzene-d6, 600
3
3
MHz): δ 8.42 (d, JHH = 7.2 Hz, 2 H, o-BC6H5), 7.55 (t, JHH = 7.2
Hz, 2 H, m-BC6H5), 7.50 (d, 3JHH = 8.4 Hz, 2 H, o-C6H5OMe), 7.40
3
3
(t, JHH = 7.2 Hz, 1 H, p-BC6H5), 6.74 (d, JHH = 8.4 Hz, 2 H, m-
C6H5OMe), 6.57−6.54 (2 H, m-C6H2Me3 + N2C3H2Mes), 6.43 (s, 1
H, m-C6H2Me3), 5.94 (d, 3JHH = 1.2 Hz, 1 H, N2C3H2Mes), 4.93 (d,
2JHH = 6 Hz, JSiH = 172 Hz, 1 H, SiH), 4.45 (t, JHH = 6 Hz, JSiH
=
1
2
1
180 Hz, 1 H, SiH), 3.65−3.57 (m, 3 H, CNCMe2CH2O), 3.39 (s, 3
H, C6H6OMe), 3.38 (d, 3JHH = 8.4 Hz, 1 H, CNCMe2CH2O), 2.02 (s,
3 H, p-C6H2Me3), 2.01 (s, 3 H, o-C6H2Me3), 1.96 (s, 3 H, o-
C6H2Me3), 1.17 (s, 3 H, CNCMe2CH2O), 1.16 (s, 3 H,
CNCMe2CH2O), 1.09 (s, 3 H, CNCMe2CH2O), 1.04 (s, 3 H,
ASSOCIATED CONTENT
* Supporting Information
■
S
1
3
CNCMe2CH2O), −13.24 (dd, JRhH = 18 Hz, JHH = 2 Hz, 1 H,
The Supporting Information is available free of charge on the
ACS Publications website at DOI: 10.1021/acs.inorg-
chem.8b03425.
RhH). 13C{1H} NMR (benzene-d6, 150 MHz): δ 194.23 (d, JRhC
=
1
1
76.5 Hz, 2C-N2C3H2Mes), 178.12 (d, JRhC = 60 Hz, CO), 80.11
(CNCMe2CH2O), 79.95 (CNCMe2CH2O), 68.42 (CNCMe2CH2O),
66.33 (CNCMe2 CH2 O), 54.50 (p-C6 H4 OMe), 28.01
(CNCMe2CH2O), 27.72 (CNCMe2CH2O), 27.41 (CNCMe2CH2O),
26.65 (CNCMe2CH2O), 20.70 (p-C6H2Me3), 19.12 (o-C6H2Me3),
18.74 (o-C6H2Me3). 15N NMR (benzene-d6, 61 MHz): δ −161
(CNCMe2CH2O trans to Si), −172 (CNCMe2CH2O trans to H),
−176 (N2C3H2Mes), −188 (N2C3H2Mes). 11B NMR (benzene-d6,
192 MHz): δ −9.9. IR (KBr, cm−1): 2961, 2933, 2836. 2562, 2531,
2170, 2016, 1903.
Spectra for compounds 2b−e and kinetics data (PDF)
Accession Codes
CCDC 1879286−1879287 contain the supplementary crys-
tallographic data for this paper. These data can be obtained
free of charge via www.ccdc.cam.ac.uk/data_request/cif, or by
emailing data_request@ccdc.cam.ac.uk, or by contacting The
Cambridge Crystallographic Data Centre, 12 Union Road,
Cambridge CB2 1EZ, UK; fax: +44 1223 336033.
NMR Kinetic Study. A stock solution of 1 (0.011 M) and
Si(SiMe3)4 (0.097 M) as an internal standard in benzene was
prepared. Portions of this solution were placed in an NMR tube, and
the initial concentration of [1] was verified by the integrated
spectrum. Excess PhSiH3 (to give concentrations greater than 0.1 M)
was added by syringe through the septum of the NMR tube. The
NMR tube was placed in the NMR probe, which had been preheated
AUTHOR INFORMATION
Corresponding Author
*E-mail: sadow@iastate.edu.
■
1
ORCID
and calibrated to the desired temperature for the reaction. H NMR
Aaron D. Sadow: 0000-0002-9517-1704
Notes
The authors declare no competing financial interest.
spectra were acquired at regular programmed intervals, and
concentrations were determined by integrating appropriate signals
in each spectrum. Plots of [1] vs time were fit using nonlinear least-
squares regression to an exponential decay curve based on the
equation [1] = [1]o e−kt for the analysis.
UV−vis Kinetic Experiments. The experiments consisted of the
following: pseudo-irreversible conditions, CO-saturated approach to
equilibrium, and equilibrium constant measurements.
Pseudo-Irreversible Conditions. Four UV−vis cuvettes, completely
filled with phenylsilane dissolved in benzene at four concentrations
(0.0192, 0.0238, 0.0291, 0.0348 M), were capped in the glovebox.
The temperature of the UV−vis chamber was set to 308.3 K, and the
sample was placed in the cavity and allowed to reach the preset
temperature. Approximately 20 μL of a solution of 1 (40 mM) was
added to the cuvette through a syringe. The absorption of the reaction
mixture was measured at 385.5 nm at programmed time intervals, and
[1] was calculated using Beer’s law. A plot of [1] vs time was fit to the
equation [1] = [1]oe−kt using nonlinear least-squares regression
analysis, to determine pseudo first-order rate constants. Subsequent
experiments varied [PhSiH3], a plot of kobs vs [PhSiH3] was linear,
and the slope provided the second-order rate constant k1app. This
procedure, performed over a range of temperatures from 296 to 322.3
K, provided the temperature-dependent second-order rate constants
for an Eyring plot.
ACKNOWLEDGMENTS
■
This research was supported by the U.S. Department of
Energy, Office of Basic Energy Sciences, Division of Chemical
Sciences, Geosciences, and Biosciences. The Ames Laboratory
is operated for the U.S. Department of Energy by Iowa State
University under contract no. DE-AC02-07CH11358.
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