Understanding the Relationship between Kinetics and Thermodynamics in CO2 Hydrogenation Catalysis

Article abstract of DOI:10.1021/acscatal.7b01673

Catalysts that are able to reduce carbon dioxide under mild conditions are highly sought after for storage of renewable energy in the form of a chemical fuel. This study describes a systematic kinetic and thermodynamic study of a series of cobalt and rhodium bis(diphosphine) complexes that are capable of hydrogenating carbon dioxide to formate under ambient temperature and pressure. Catalytic CO₂ hydrogenation was studied under 1.8 and 20 atm of pressure (1:1 CO₂/H₂) at room temperature in tetrahydrofuran with turnover frequencies (TOF) ranging from 20 to 74000 h^-1. The catalysis was followed by ¹H and ³¹P NMR spectroscopy in real time under all conditions to yield information about the rate-determining step. The cobalt catalysts displayed a rate-determining step of hydride transfer to CO₂, while the hydrogen addition and/or deprotonation steps were rate limiting for the rhodium catalysts. Thermodynamic analysis of the complexes provided information on the driving force for each step of catalysis in terms of the catalyst hydricity (G°_H), acidity (pK_a), and free energy for H₂ addition (G°_H2). Linear free-energy relationships were identified that link the kinetic activity for catalytic hydrogenation of CO₂ to formate with the thermodynamic driving force for the rate-limiting steps of catalysis. The catalyst exhibiting the highest activity, Co(dmpe)₂H, was found to have hydride transfer and hydrogen addition steps that were each downhill by approximately 6 to 7 kcal mol^-1, and the deprotonation step was thermoneutral. This indicates the fastest catalysts are the ones that most efficiently balance the free energy relationships of every step in the catalytic cycle.

Full text of DOI:10.1021/acscatal.7b01673

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Article
Understanding the Relationship Between Kinetics and
Thermodynamics in CO2 Hydrogenation Catalysis
Matthew S Jeletic, Elliott Hulley, Monte L Helm, Michael T. Mock,
Aaron M. Appel, Eric S. Wiedner, and John Charles Linehan
ACS Catal., Just Accepted Manuscript • DOI: 10.1021/acscatal.7b01673 • Publication Date (Web): 25 Jul 2017
Downloaded from http://pubs.acs.org on July 26, 2017
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Understanding the Relationship Between Kinetics and Thermodynamics in CO2
Hydrogenation Catalysis
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Authors: Matthew S. Jeletic, Elliott B. Hulley, Monte L. Helm, Michael T. Mock, Aaron M. Appel,
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Eric S. Wiedner,* and John C. Linehan*
Catalysis Science Group, Pacific Northwest National Laboratory, Richland, Washington, 99352, USA
* E-mail: eric.wiedner@pnnl.gov, john.linehan@pnnl.gov
Abstract
Catalysts that are able to reduce carbon dioxide under mild conditions are highly sought after for
storage of renewable energy in the form of a chemical fuel. This study describes a systematic kinetic and
thermodynamic study of a series of cobalt and rhodium bis(diphosphine) complexes that are capable of
hydrogenating carbon dioxide to formate under ambient temperature and pressure. Catalytic CO2
hydrogenation was studied under 1.8 and 20 atm of pressure (1:1 CO2:H2) at room temperature in
tetrahydrofuran with turnover frequencies (TOF) ranging from 20 to 74,000 h-1. The catalysis was
followed by 1H and 31P NMR spectroscopy in real time under all conditions to yield information about the
rate determining step. The cobalt catalysts displayed a rate determining step of hydride transfer to CO2,
while the hydrogen addition and/or deprotonation steps were rate limiting for the rhodium catalysts.
Thermodynamic analysis of the complexes provided information on the driving force for each step of
catalysis in terms of the catalyst hydricity (G°H-), acidity (pKa), and free energy for H2 addition (G°H2).
Linear free-energy relationships were identified that link the kinetic activity for catalytic hydrogenation
of CO2 to formate with the thermodynamic driving force for the rate-limiting steps of catalysis. The
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catalyst exhibiting the highest activity, Co(dmpe)2H, was found to have hydride transfer and hydrogen
addition steps that were each downhill by approximately 6 to 7 kcal mol–1, and the deprotonation step was
thermoneutral. This indicates the fastest catalysts are the ones that most efficiently balance the free energy
relationships of every step in the catalytic cycle.
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Keywords
Carbon dioxide; hydrogenation; hydricity; homogeneous catalysis; cobalt; rhodium
Introduction
Photocatalytic and electrocatalytic reduction of CO2 to liquid fuels is widely pursued as a means
to store energy generated from geographically disperse renewable sources, such as wind and solar.1-6 As
an alternative to direct reduction using protons and electrons, CO2 can also be hydrogenated using clean
H2 obtained from renewable sources.7-11 Molecular catalysts for hydrogenation of CO2 are selective for
production of formic acid / formate, an attractive chemical hydrogen storage material.12-14 The noble
metals, specifically Ir, Rh and Ru, are often the focus for catalytic hydrogenation of CO2 to formic acid.15-
34 Catalysts based on the more abundant metals Ni, Co, Fe, and Cu have been described,35-51 though these
catalysts often do not rival the activities of noble metal catalysts under optimized conditions. A common
link in most of these optimized catalyst systems is the requirement for high temperatures and pressures,52-
53 which is undesirable for decentralized conversion of CO2 using dispersed renewable energy sources.1,
54-55
For example, Nozaki’s Ir-pincer complex is inactive at 25 °C (>40 atm), and yields a low turnover
frequency (TOF) of 1 h-1 at 100 °C and 1 atm.25-26 Fujita and Himeda et al., demonstrated catalysis under
ambient conditions using Ir-bipyridyl catalysts, though the TOF’s were low (<100 h-1) compared to
catalysts operating at higher pressures and temperature.22-24 We previously reported that Co(dmpe)2H
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(dmpe = 1,2-bis(dimethylphosphino)ethane) is much more active, although with the need for a very strong
base, than any other catalyst system under ambient conditions (25 °C and 1 atm) and quite competitive as
one of the fastest systems to date at higher pressures.43-44
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Rational design of improved catalysts for hydrogenation of CO2 under ambient conditions requires
an understanding of the mechanism. Most CO2 hydrogenation catalysts are proposed to follow a basic
three step mechanism: (1) hydrogen addition to the metal complex, (2) deprotonation of the H2-adduct by
exogeneous base, and (3) hydride transfer to CO2 to afford formate (Scheme 1).22-25, 38-39, 41, 56-57 For some
catalysts, the non-innocent ligand scaffold can play a role by storing hydride equivalents,25-26, 58 though
this does not change the overall nature of the three mechanistic steps of CO2 hydrogenation. A key factor
for achieving high catalytic activity with Co(dmpe)2H under ambient conditions is to use Verkade’s base
+
(Vkd), which is basic enough (pKa(VkdH+) = 33.6 in acetonitrile)59 to deprotonate Co(dmpe)2H2 (pKa =
33.7).60 When using Vkd as the base, hydride transfer from Co(dmpe)2H to CO2 becomes rate-limiting for
catalysis, despite having a favorable reaction free energy of –8 kcal mol–1.43
Scheme 1. General catalytic cycle for CO2 hydrogenation.
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In this study, we sought to learn how to optimize the thermodynamic driving force of each catalytic
step in order to maximize the catalytic turnover frequency. By examining a series of cobalt and rhodium
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bis(diphosphine) complexes, M(P2)2 , we demonstrate that the kinetics of catalytic CO2 hydrogenation are
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strongly influenced by the thermodynamic properties of the catalyst. The driving force of the catalytic
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+
steps are determined by the free energy for H2 addition to M(P2)2 (ΔG°H2), the pKa of M(P2)2H2 , and the
hydride donor ability (or hydricity) of M(P2)2H (ΔG°H–). Our results reveal the presence of a linear free
energy relationship in which the catalytic TOF is correlated with the thermodynamic driving force of the
rate-limiting step of catalysis. Importantly, these findings provide critical insight into the properties that
make a catalyst active for hydrogenation of CO2 under ambient conditions.
Results
Catalytic Hydrogenation of CO2. The cobalt and rhodium complexes that were studied for catalytic CO2
hydrogenation in this work are displayed in Chart 1. Each of these complexes has been previously
reported, or was synthesized by extension of literature procedures (see the Experimental section for more
details). For cobalt, the Co(P2)2H form of the catalyst was used. In general, the Rh(P2)2H complexes are
much less stable than their Co counterparts and could not be isolated cleanly. Rh(dppe)2H is an
exception,61 and its increased stability most likely results from increased steric protection imparted by the
phenyl rings and the decreased electron density at the rhodium center. We previously demonstrated that
+
Co(P2)2 is an important catalytic intermediate and there is no difference in TOF between it and
+
Co(P2)2H.44 By extension, we expect that Rh(P2)2 is suitable for catalytic studies instead of the unstable
Rh(P2)2H complexes.
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Chart 1. a) Cobalt and rhodium complexes studied for catalytic CO2 hydrogenation. b) Structures of the
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diphosphine ligands.
Catalytic hydrogenation of 1:1 CO2:H2 at was studied at both 1.8 atm and 20 atm using an excess
of Verkade’s Base (Vkd, >400 mM) and different cobalt and rhodium catalysts at room temperature in
THF-d8. Figure 1 shows a representative plot of turnovers (mol of [VkdH]+[HCO2]– per mol of catalyst)
versus time for Co(BPE5)2H at 20 atm of 1:1 CO2:H2. Pseudo-first order catalytic rate constants were
measured from the linear portion of the time-dependent kinetic profiles. These catalytic rate constants are
equal to the catalytic turnover frequency (TOF), i.e. moles of [VkdH]+[HCO2]– per mole of catalyst per
hour, under the specified conditions and are listed in Tables 1-2. At higher concentrations of the catalysts,
the rate of catalysis was mass-transport limited, specifically by the rate of the gas-liquid mixing. Therefore
the catalyst concentration was decreased until the TOF remained similar at two different catalyst
concentrations (Tables S2-S3), indicating that the reaction was no longer mass-transport limited. Several
of the catalysts were too slow at 1.8 atm to measure a significant amount of [VkdH]+[HCO2]– within the
first hour, so the TOF of these catalysts were not measured under these conditions.
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Figure 1. A plot of turnovers (mol [VkdH]+[HCO2]– / mol catalyst) versus time for 0.036 mM
Co(BPE5)2H at 20 atm of 1:1 H2:CO2 and 570 mM Vkd in THF-d8 at 21 °C. The reaction ceases when all
of the Vkd has been consumed. The line is the best linear fit and was used to determine the TOF (R2 =
0.99).
Table 1. Catalytic conversion of CO2 and H2 to formate with Verkade’s base at 1.8 atm.a
Initial Loading
Catalyst
TOF (h-1) b Turnovers b
(mM)
0.28
2.6
Co(dmpe)2H
Co(BPE5)2H
Co(dmpbz)2H
Co(depe)2H
6400
540
1900 c
170
200
20
2.6
270
25
21
+
Rh(dmpbz)2
0.27
0.30
2.7
3100 d
1600 d
90 d
1100
280
55
+
Rh(depe)2
+
Rh(dmpe)2
+
Rh(dmpe)2
2.7
640 e
88 e
a Catalytic conditions; 500µL THF-d8, 1.8 atm 1:1 CO2:H2, 21 °C, 400-430 mM 2,8,9-triisopropyl-2,5,8,9-tetraaza-1-
phosphabicyclo[3,3,3] undecane (Verkade’s base), typical run time 2 h or less. b Uncertainties are 10%. c Previously reported.
t
d Initial TOF. e tert-Butylimino-tris(dimethylamino)phosphorane, pKa ~ 38 for P4 BuH+ in MeCN,62 was used instead of
Verkade’s base.
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Table 2. Catalytic conversion of CO2 and H2 to formate with Verkade’s base at 20 atm. a
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Initial Loading
Catalyst
TOF (h-1) b Turnovers b
(mM)
0.040
Co(dmpe)2H
Co(BPE5)2H
Co(dmpbz)2H
Co(depe)2H
Co(dedpe)2H
Co(dppe)2H
74000 c
13000
11000
610
9400 c
9700
6700
1300
26
9
0.036
0.037
0.36
22
8.3 d
0.038
0.43
0.38
7.1
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11 e
17 e
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Rh(dmpbz)2
11000 f
3000 f
7100
1000
100
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Rh(depe)2
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f
Rh(dmpe)2
145
Rh(dppe)2H
130
33 e
70
93 e
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Rh(dcpe)2
2.6 d
aCatalytic conditions; 350µL THF-d8, 20 atm 1:1 CO2:H2, 21 °C, 470-600 mM 2,8,9-triisopropyl-2,5,8,9-tetraaza-1-
phosphabicyclo[3,3,3] undecane (Verkade’s base), typical run time 2 h or less. b Uncertainties are 10%. c Previously reported.
d solubility limited catalyst loading. e 40 atm. f initial TOF.
The time-resolved 1H and 31P{1H} NMR spectra reveal the presence of only Co(P2)2H species over
the course of the reaction at both low and high pressure, suggesting a rate determining step of hydride
transfer from Co(P2)2H to CO2 in all cases. The time-dependent kinetic profiles of the cobalt-catalyzed
reactions remain linear over the course of the reaction, though an induction period of ≤ 10 minutes was
observed for some of the cobalt catalysts. Time-resolved NMR spectroscopy experiments did not reveal
any other species other than Co(P2)2H complex during the induction period, and visual inspection of the
reaction did not reveal any noticeable color changes or precipitate formation. Gas-liquid mixing does not
1
appear to be the source of the induction period, as the H NMR resonance for dissolved H2 remained
constant during the induction period.
+
At low concentrations of the Rh(P2)2 catalysts, the kinetic plots display a significant decrease in
catalyst activity over time, with catalytic activity ceasing after about two hours. We were unable to monitor
the catalyst when suspected decomposition is occurring due to the low catalyst loadings. At higher
+
loadings where the Rh(P2)2 catalyst could be monitored, the maximum number of turnovers is reached
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before decomposition occurs. Mass-transport limitations complicated the determination of an accurate
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TOF at high loadings of Rh(P2)2 , so the initial TOFs measured at low concentrations of the Rh(P2)2
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catalysts are reported in Tables 1-2.
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Rh(dmpe)2H2 is the only species observed in time-resolved NMR spectra at 20 atm of 1:1 CO2:H2,
indicating that deprotonation is the rate-determining step under these conditions. At 1.8 atm of 1:1
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CO2:H2, time-resolved NMR spectra reveal the presence of Rh(dmpe)2 and Rh(dmpe)2H2 in about equal
concentrations. The lowering of pressure directly influences the kinetic barrier for hydrogen addition to
make it competitive with deprotonation such that two concurrent steady-state intermediates exist. In
+
1
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contrast, only Rh(depe)2 is observed in time-resolved H and P NMR spectra at both 1.8 atm and 20
atm, supporting a hydrogen addition rate determining step for this complex.
While the focus for this study is more on TOF, the catalysts also afford a modest number of
turnovers. Importantly, the number of turnovers is not limited by catalyst decomposition (for the cobalt
catalysts), but instead by the amount of Vkd present. Diluting the catalyst concentration infinitely is not a
viable solution for increasing the number of turnovers, as residual impurities in the reagents
disproportionately start to affect catalyst performance at lower concentrations of catalyst.
Since the rhodium catalysts exhibit a different resting state than the cobalt catalysts, we studied
+
the mechanism of Rh(dmpbz)2 in more detail to verify that the rhodium and cobalt catalysts operate by
similar mechanisms. [Rh(dmpbz)2](O3SCF3) displayed low solubility in tetrahydrofuran, so one
equivalent of NaBArF (ArF = 3,5-(CF3)2C6H3) was added to solubilize the complex. 31P{1H} and 1H NMR
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+
+
spectra indicate that Rh(dmpbz)2 is in equilibrium with the dihydride complex, cis-Rh(dmpbz)2H2 (56%
+
conversion), under 1.8 atm of H2. Diagnostic broad peaks attributable to cis-Rh(dmpbz)2H2 resonate at
35.0 and 33.8 ppm in the 31P{1H} NMR spectrum and –9.28 (hydride, d, 1JHRh ~ 115 Hz) ppm in the 1H
+
NMR spectrum. Only Rh(dmpbz)2 and Rh(dmpbz)2H (broad resonances at –10.80 (hydride) and 21.6
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ppm (JPRh = 141 Hz) in the H and P{1H} NMR spectra, respectively) were observed when this
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experiment was repeated in the presence of ten equivalents of Vkd base. Addition of 1.8 atm of CO2
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1
converts Rh(dmpbz)2H back into Rh(dmpbz)2 as observed in the H and 31P{1H} NMR spectra. In a
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separate experiment, pressurizing Rh(dmpbz)2 with 6.7 atm of H2 shifts the equilibrium to favor only cis-
+
Rh(dmpbz)2H2 . Based on these results, we are reasonably certain that the rhodium complexes follow the
same mechanism as the cobalt catalysts, but have a different rate determining step. A similar mechanism
was concluded recently in a computational study of Rh(P2)2 complexes.63 As a last note, unlike the
+
+
+
Co(P2)2 catalysts, the Rh(P2)2 complexes do not react to form unwanted side products (non-catalytic
dimer complex or CO complex).44
Thermodynamics
Accurate relative thermodynamic values of the different complexes are needed to properly assess
the correlations between driving force and catalytic turnover frequency. Previous studies have
demonstrated the challenges associated with experimental measurement of hydricity values for strongly
hydridic metal hydrides, such as Co(dmpe)2H.64 Computational methods can often provide accurate
thermodynamic values for transition metal hydrides, but complexes possessing ligands with large
conformational degrees of freedom (such as depe) still pose a significant challenge.60, 65-66 As a result, we
estimated the relative hydricities of the Co(P2)2H and Rh(P2)2H complexes through empirical correlation
(described below) instead of by direct experimental measurement or computation.
Two of the Co(P2)2H complexes used in this study have known hydricity values: Co(dppe)2H
(ΔG°H– = 49.9 kcal mol–1)67-68 and Co(dmpe)2H (ΔG°H– ≈ 36.3 kcal mol–1).65, 69 We relied on two
considerations to estimate hydricities for the remaining Co(P2)2H complexes. First, Co(P2)2H and
Ni(P2)2H+ complexes are isoelectronic. Second, the hydricities of Ni(P2)2H+ complexes are linearly
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correlated with the Ni(II/I) couples of the parent Ni(P2)22+ complexes.70 The analogous reduction potential
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for Co(P2)2 is the Co(I/0) couple (a d8 to d9 reduction), which has been previously reported for
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Co(dppe)2 and Co(dedpe)2 in acetonitrile.67, 69 Cyclic voltammetry experiments were performed in
acetonitrile to measure the potential of the Co(I/0) couple for the remaining Co(P2)22+ complexes, except
for Co(BPE5)22+ which was insoluble in acetonitrile and other nitrile solvents. A plot of hydricity versus
reduction potential was constructed using Co(dppe)2H and Co(dmpe)2H (both having known hydricities),
as well as the the Ni(P2)2H+ analogs of the Co(P2)2H complexes in this study (see the SI for more details).
From the linear fit to this plot, hydricity values were estimated for the remaining Co(P2)2H complexes
from their Co(I/0) reduction potentials (Table 3). It is important to note that the Co(I/0) couples of
2+
2+
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+
+
Co(dmpbz)2 and Co(dmpe)2 were chemically irreversible in acetonitrile. As a result, the hydricity values
may have low absolute accuracy, however, the relative trend in hydricities is expected to be accurate.
Table 3. Summary of Estimated Thermochemical Values.a
G°H– MI(H)
(kcal mol–1)
49.9 b
G°H2 MI
(kcal mol–1)
-5.0 b
Complex
pKa MIII(H)2
+
Co(dppe)2
22.8 b
27.4
+
Co(dedpe)2
44.2
40.6
-5.5
+
Co(depe)2
30.3
-5.9
+
Co(dmpbz)2
38.0
36.3 c
32.3
33.7 c
-6.1
+
Co(dmpe)2
-6.3 c
4.2 d
3.8 d
0.9 e
0.2 f
-0.4 e
+
Rh(dcpe)2
39.9
24.0
+
Rh(dppe)2
36.2
28.3 e
26.4
34.3 e
+
Rh(depe)2
+
Rh(dmpbz)2
27.4
26.6 e
35.5
36.5 e
+
Rh(dmpe)2
a Values are estimated as described in the text unless noted otherwise. b Experimental value, see Ref 67-68
c Computational value, see Ref 60, 65. d Experimental value, see Ref 71. e Experimental value, see Ref 64. f
Experimental value, this work.
.
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Hydricity values have been previously measured for Rh(dmpe)2H (ΔG°H– = 26.6 kcal mol–1) and
Rh(depe)2H (ΔG°H– = 28.3 kcal mol–1),64, 68 as well as for the dihydride counterparts Rh(dmpe)2H2
+
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9
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(ΔG°H– = 49.9 kcal mol–1) and Rh(depe)2H2 (ΔG°H– = 51.9 kcal mol–1).68, 71 These values correspond to
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a 24 kcal mol–1 difference between the Rh(P2)2H and Rh(P2)2H2 complexes of the same diphosphine
ligand. A similar difference in hydricities (25 kcal mol–1) was measured for Rh(depx)2H and Rh(depx)2H2
+
(depx = α,αʹ-bis(diethylphosphino)xylene), complexes that contain a diphosphine with a large bite angle.72
Therefore we estimated the hydricities of Rh(dcpe)2H and Rh(dppe)2H by subtracting 24 kcal mol–1 from
+
+ 68, 71
+
the experimental hydricities of Rh(dcpe)2H2 and Rh(dppe)2H2 .
The pKa values of the Rh(P2)2H2
complexes were determined from a thermochemical cycle using the hydricity of Rh(P2)2H, the free energy
for H2 addition to Rh(P2)2 , and the constant for heterolytic cleavage of H2 in acetonitrile.68
+ 71
Examination of the thermochemical data indicates that the hydricity of Rh(P2)2H is inversely
+
+
linearly correlated with the pKa of Rh(P2)2H2 and with the free energy for H2 addition to Rh(P2)2 (see
+
SI). Thus as Rh(P2)2H becomes a stronger hydride donor, Rh(P2)2H2 becomes harder to deprotonate and
+
H2 addition to Rh(P2)2 becomes more favorable. These findings were used to estimate additional
+
thermochemical data for the complexes in this study. First, the free energy for H2 addition to Rh(dmpbz)2
was measured to be 0.2 kcal mol–1 from the reactivity studies described in the previous section. This value
+
was used to estimate the hydricity of Rh(dmpbz)2H and the pKa of Rh(dmpbz)2H2 using the linear
+
+
relationships of the other rhodium complexes. Second, pKa values of Co(dmpbz)2H2 , Co(depe)2H2 , and
+
Co(dedpe)2H2 were estimated by assuming a linear relationship between the pKa and the hydricity of
Co(P2)2H, anchored to the pKa values of Co(dppe)2H2+ 67 and Co(dmpe)2H2+.60, 69 The free energy for H2
+
addition to Co(P2)2 was then determined from a thermochemical cycle using the hydricity of Co(P2)2H,
the pKa of Co(P2)2H2 , and the constant for heterolytic cleavage of H2 in acetonitrile.68 Full details of these
+
estimations are provided in the SI.
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In this study, we sought to gain a better understanding of how to balance the thermodynamic
driving force for each step of catalytic hydrogenation of CO2 in order to maximize the catalytic turnover
frequency. This objective relies on being able to correlate the thermodynamic properties of the catalysts
with the kinetic data for catalytic CO2 hydrogenation. However, the thermodynamic properties were
determined for acetonitrile as a solvent, whereas the kinetic data were collected in tetrahydrofuran. Leito
and coworkers have demonstrated that pKa values of organic acids are consistently 5-6 units higher in
acetonitrile than in tetrahydrofuran.73-74 Therefore we assume that the thermodynamic parameters of the
catalysts have the same relative values between species in both acetonitrile and tetrahydrofuran, thereby
allowing a meaningful comparison of driving force and reaction rate.
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The rate determining step for the Co(P2)2 catalysts was found to be hydride transfer to CO2
(equation 1). Therefore, the catalytic turnover frequencies are equal to the pseudo-first order rate constant
for hydride transfer at a given pressure of CO2. A plot of the catalytic TOF (on a log scale) versus the
hydricity of Co(P2)2H is shown in Figure 2 for CO2 partial pressures of 0.9 atm and 10 atm (total pressure
of 1.8 atm and 20 atm, respectively). Notably, the hydricity and catalytic rate are linearly correlated at
each pressure. This finding is consistent with a recent computational study on Fe and Co complexes
containing tripodal ligands, which predicted a linear correlation between hydricity and the activation
barrier for hydride transfer to CO2.75 To the best of our knowledge, our present results are the first
experimental demonstration of a linear free energy relationship for hydride transfer to CO2 in a series of
isostructural transition metal hydride complexes.
+
–
Co(P2)2H + CO2 ⇌ Co(P2)2 + HCO2
(1)
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Figure 2. Plot of the logarithm of the catalytic turnover frequency versus the hydricity of Co(P2)2H. The
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data contained in this plot are tabulated in the SI.
Analysis of the linear free energy correlation according to the Bell-Evans-Polyani principle76
affords equation 2, where α describes the position of the transition state along the reaction coordinate
(0 ≤ α ≤ 1) and C is a constant (see the SI). At 0.9 atm of CO2, the slope of the free energy correlation
corresponds to α = 0.78, indicating a late transition state that structurally resembles the product, formate,
as opposed to the reactant, CO2. Previous computational studies of Co(dmpe)2H identified two kinetically
competent mechanisms for hydride transfer: direct outer-sphere hydride transfer to CO2, and an
associative mechanism where CO2 first binds to Co before formation of the C–H bond (Figure 3).77 Both
mechanisms require a rehybridization of CO2 from linear to bent in the rate-determining transition state
(RDTS), consistent with the experimental determination of a late transition state for hydride transfer.
log(TOF) = –(α/1.364)×ΔG°H– + C
(2)
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Figure 3. Previously calculated mechanisms for hydride transfer from Co(dmpe)2H to CO2.
Few studies examine both the thermodynamics and kinetics for H– transfer to CO2, so it is difficult
to compare the activity for hydride transfer from Co(P2)2H to other catalysts. An exception is
Ru(terpy)(bpy)H+, for which both thermodynamic and kinetic data is available for hydride transfer in
acetonitrile. The hydricity of Ru(terpy)(bpy)H+ is 39.3 kcal mol–1 in acetonitrile,68, 78 corresponding to a
favorable driving force of ~5 kcal mol–1 for hydride transfer to 1 atm of CO2. A second-order rate constant
of 0.018 M–1 s–1 was reported for hydride transfer from Ru(terpy)(bpy)H+ to CO2 in acetonitrile at 25 °C.79
This second-order rate constant is equivalent to a pseudo first-order rate constant (TOF) of 19 h–1 at 1 atm
of CO2, and is much smaller than the TOF of 150 h–1 predicted from the hydricity of Ru(terpy)(bpy)H+
using the free energy correlation of Figure 2. Therefore hydride transfer to CO2 is intrinsically slower for
Ru(terpy)(bpy)H+ than for Co(P2)2H, possibly due to a difference in relative stabilities of the metal
+
complexes resulting from hydride transfer. Four-coordinate Co(P2)2 (16 e–, d8) should be a more stable
species than five-coordinate Ru(terpy)(bpy)2+ (16 e–, d6).
The BPE5 ligand was considered based on previous evidence by Field that dmpe and BPE5 should
have similar chelation strength.80 We were unable to estimate thermodynamic properties for Co(BPE5)2
+
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due to its insolubility in nitrile solvents. However, the catalytic turnover frequency of Co(BPE5)2H was
found to lie between Co(dmpe)2H and Co(dmpbz)2H, suggesting that the hydricity of Co(BPE5)2H is
between that of Co(dmpe)2H (36.3 kcal mol–1) and Co(dmpbz)2H (38.0 kcal mol–1).
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+
The Rh(dmpe)2 catalyst has a low average TOF of 200 h-1 at 20 atm due to a large mismatch
between the pKa of Rh(dmpe)2H2 and Vkd (approximately 3 units).43-44 The mismatch is also apparent
+
when examining the time resolved multinuclear NMR data, as the rate determining step is clearly
+
deprotonation since the only species observed during catalysis is Rh(dmpe)2H2 . Notably, the TOF
increases from 90 h-1 to 640 h-1 when a stronger base, tert-Butylimino-tris(dimethylamino)phosphorane,
N′-tert-Butyl-N,N,N′,N′,N″,N″-hexamethylphosphorimidic triamide (P4 Bu) (pKa ~ 38 in CH3CN)),62 is
t
+
+
used at 1.8 atm. In contrast, the other Rh(P2)H2 species have a lower pKa than Rh(dmpe)2H2 , and the
rate determining step changes from deprotonation to hydrogen addition. Therefore, the free energy for H2
+
addition to Rh(P2)2 (equation 3) is the appropriate thermodynamic value to compare to the catalytic TOF
+
of the rhodium catalysts (except for Rh(dmpe)2 ). A plot of the catalytic TOF (on a log scale) versus the
free energy for H2 addition is shown in Figure 4 for catalysis at an H2 partial pressure of 10 atm (total
pressure of 20 atm), resulting in a linear correlation between log(TOF) and the driving force.
+
+
Rh(P2)2 + H2 ⇌ Rh(P2)2H2
(3)
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Figure 4. Plot of the logarithm of the catalytic turnover frequency at 20 atm of 1:1 H2:CO2 versus the
free energy for H2 addition to Rh(P2)2 . The data contained in this plot is tabulated in the SI.
+
Using the comprehensive thermodynamic information for the different cobalt and rhodium
catalysts, we can analyze the factors that promote catalysis under low pressure, room temperature
conditions. A reaction coordinate diagram showing the relative free energy values for each elementary
catalytic step of six catalysts that operate under ambient conditions (1.8 atm of total pressure, 25 °C) is
shown in Figure 5. The net reaction free energy is defined by the identity of the substrates (H2, CO2, Vkd)
–
and products (HCO2 , VkdH+), while the relative energies of the catalytic intermediates are defined by the
thermodynamic properties of each catalyst. An important point for catalyst design is that the sum of the
driving forces for each catalytic step must equal the net reaction free energy. As a result, a catalyst that
possesses a large driving force for a single step will necessarily have a low driving force for one or more
other steps in the catalytic cycle.
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Figure 5. Reaction coordinate diagram showing the relative free energies of six catalysts that operate
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under ambient conditions (1.8 atm, 25 °C). The lines are not correlations and are only added to guide the
eye.
+
Co(dmpe)2 is the fastest of the catalysts that are active under ambient conditions. As seen in
Figure 5, the hydride transfer and H2 addition steps are both downhill in energy, while the deprotonation
+
+
step is thermoneutral. The other cobalt catalysts, Co(dmpbz)2 and Co(depe)2 , have a greater driving
+
force for deprotonation than Co(dmpe)2 , but less driving force for hydride transfer to CO2. The extra
driving force for deprotonation is “wasted” energy since this is not the rate determining step; as a result,
+
+
+
Co(dmpbz)2 and Co(depe)2 are slower catalysts than Co(dmpe)2 . In principal, catalysis could be
+
improved over Co(dmpe)2 by increasing the driving force for hydride transfer, decreasing the driving
force for H2 addition, and maintaining the thermoneutrality of deprotonation. As an extreme example,
+
+
these conditions are fulfilled in Rh(dmpbz)2 and Rh(depe)2 , which have a very strong driving force for
hydride transfer and little to no driving force for H2 addition or deprotonation. However, the
+
thermodynamics for rhodium catalysts have shifted too far, and the catalysts are slower than Co(dmpe)2
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due to a change in the rate determining step from hydride transfer to H2 addition. Thus, the rhodium
catalysts possess too much driving force for the hydride transfer step, and not enough driving force in the
H2 addition step. In a computational study of an Ir(PNP) hydride catalyst, Hazari and Bernskoetter
similarly concluded that hydride transfer to CO2 was too favorable and H2 addition was not favorable
enough.81
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These results demonstrate that in a family of structurally related catalysts, the fastest catalyst is the
one that most efficiently balances the linear free energy correlations of all the steps in the catalytic cycle
(in this case, hydride transfer, H2 addition, and deprotonation). As a result, it is unlikely that a given
catalyst will display remarkable activity for CO2 hydrogenation under ambient conditions if any of the
reaction steps are energetically unfavorable by more than 1-2 kcal mol–1. This requires avoiding excess
driving force for any one step (i.e. greater than the driving force of the net reaction), since the
thermodynamic properties of the catalyst (ΔG°H–, ΔG°H2, and pKa) are fundamentally connected in a
thermochemical cycle.
The optimal balance in driving force between the catalytic steps will likely vary between different
catalyst families that have varying correlations between driving force and rate for each step of catalysis.
An illustrative example is the complex [Cp*Ir(OH)2(6,6ʹ-dihydroxy-2,2ʹ-bipyridyl)]2+, which is also an
active catalyst for hydrogenation of CO2 at room temperature and pressure.23 Computational studies on
this Ir complex suggest a similar reaction pathway of H2 addition, deprotonation, and hydride transfer. Of
these three steps, H2 addition was found be the rate-limiting step for catalysis despite having a stronger
driving force (–2.4 kcal mol–1) than deprotonation (+2.5 kcal mol–1) or hydride transfer (–0.1 kcal mol–1).
An initial TOF of 27 h–1 was reported for this Ir catalyst at 1 atm of 1:1 H2:CO2 in aqueous bicarbonate
solution, which would correspond to TOF of 49 h–1 at 1.8 atm of 1:1 H2:CO2 (the pressure used in the
+
present study). Rh(dmpbz)2 displays a larger TOF of 3100 h–1 at 1.8 atm of 1:1 H2:CO2 despite having a
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lower driving for H2 addition (+0.2 kcal mol–1) than the Ir complex. This difference in TOF is likely due
in part to the greater solubility of H2 in organic solvents than in water.82
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Summary and Conclusions
While many homogeneous catalysts for CO2 hydrogenation to formate have been reported, most
require temperatures and pressures in excess of 80 °C and 20 atm. Several groups recognize the need to
design catalysts that operate at low temperature and pressure, but very few catalysts have been reported
that are active under ambient conditions. We examined the kinetics and thermodynamics for CO2
+
hydrogenation in a series of Co(P2)2H and Rh(P2)2 complexes in order to gain a better understanding of
the factors that lead to fast catalysis under ambient conditions. Hydride transfer was rate limiting for the
+
cobalt catalysts, while hydrogen addition was rate limiting for all rhodium catalysts except Rh(dmpe)2 .
Both the cobalt and rhodium complexes displayed a linear free energy correlation between the catalytic
turnover frequency and the thermodynamic driving force for the rate limiting steps. The fastest catalyst in
this study was Co(dmpe)2H (TOF = 6400 h–1 at 1.8 atm), which had a similar driving force for both hydride
transfer (–7.6 kcal mol–1) and H2 addition (–6.2 kcal mol–1), with the deprotonation step being
thermoneutral. This work demonstrates that catalyst design can be greatly aided by consideration of the
thermodynamics of every step in the catalytic cycle. Importantly, the catalytic activity is maximized when
the linear free energy relationships of the fundamental steps are properly balanced.
Experimental Section
General Considerations: All manipulations were performed under an N2 atmosphere unless stated
otherwise. Tetrahydrofuran, acetonitrile, ether, and hexane were purified by sparging with N2 and passage
through neutral alumina using an Innovative Technology, Inc., PureSolv solvent purification system.
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Triethylamine, DBU, the phosphazene base and Verkade’s base were purchased from Sigma-Aldrich and
used without further purification. Triethylamine and DBU were dried over sodium and 4Å sieves until a
Karl-Fisher reading of less than 20 ppm was obtained. THF-d8 and CD3CN were ordered from Cambridge
Isotopes. THF-d8 was dried over a sodium-potassium ball, and CD3CN was dried over P2O5. CoCl2,
[Rh(COD)2](CF3SO3), dmpe, depe and dmpm were purchased from Strem Chemical and used without
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further purification. Syntheses of BPE5,83 dmpbz,84 Co(dmpe)2H,69, 85 Co(dedpe)2H,69 Co(depe)2H,86
87
[Rh(dmpe)2](CF3SO3)/Rh(dmpe)2H,64,
[Rh(depe)2](CF3SO3),71
[Rh(dcpe)2](CF3SO3),71
[Rh(dppe)2](CF3SO3),71 [Co(CH3CN)6](BF4)2,88 and Rh(dppe)2H61 were previously reported elsewhere.
Gases were purchased from Oxarc and were delivered to the NMR tube directly from the tank
through a gas manifold line at pressures less than 1.8 atm and by an ISCO pump at pressures greater than
6.7 atm. The PEEK tubes were designed and built at Pacific Northwest National Laboratory.89-90 The
1
glass capillary standard was prepared by adding Co(NO3)2 to water until the resonance in the H NMR
spectrum appeared at a desirable location. All gas mixtures were measured by gas chromatography. GC
analyses were carried out using a 100/120 Carbosieve S-II stainless steel column with a length of 10 feet
and a diameter of 1/8 in. purchased from Supelco. A thermal conductivity detector was used, and the oven
temperature was ramped from 100 to 175 °C. Argon was used as the carrier with a flow rate of 15 mL/min.
A CH Instruments 660C potentiostat was used for all voltammetry experiments. Experiments were
performed in a nitrogen glovebox (23 ± 2 °C) using a standard three-electrode configuration. A 1 mm
diameter PEEK-encased glassy carbon disc (ALS) was used as the working electrode, a glassy carbon rod
(Structure Probe, Inc.) was used as the counterelectrode, and a bare silver wire was used as the
pseudoreference electrode. Ferrocene was used as an internal standard, and potentials are referenced to
the Cp2Fe+/0 couple.
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Representative Synthetic Procedures:
Trans-[Co(dmpbz)2(CH3CN)2](BF4)2. To a 50 mL round bottom flask was added
[Co(CH3CN)6](BF4)2 (978 mg, 2.04 mmol) and CH3CN (10 mL). To this solution was added dmpbz (829
mg, 4.18 mmol) dropwise. The solution turned from red to brown and was stirred for 10 min at room
temperature. The solution was then allowed to sit for 2 h. A brown precipitate formed, which was filtered
off. Crystals were grown by diffusing ether into the remaining reaction solution, yielding large dark purple
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1
blocks (1.19 g, 77%). Note crystals become orange upon desolvation of one CH3CN molecule. H NMR
(paramagnetic, CD3CN):  7.67 (br, 2H), 5.26 (br, 2H), 2.10 (br, CH3CN, 6H), -3.76 (br, PCH3, 12H).
Anal. Calcd for C24H38B2CoF8N2P4: %C, 40.54; %H, 5.39; %N, 3.94. Found %C, 40.61; %H, 5.20; %N,
3.96.
A
similar procedure was used to synthesize [Co(depe)2CH3CN](BF4)2 and
[Co(BPE5)2(CH3CN)x](BF4)2; see the SI for full details.
Co(dmpbz)2H. To a 50 mL round bottom flask was added [Co(dmpbz)2(CH3CN)2](BF4)2 (1.18
g, 1.66 mmol) and CH3CN (10 mL). To this solution was slowly added KC8 (475 mg, 3.52 mmol). The
solution turned from purple to red and was stirred for 20 min at room temperature. The flask was then
sealed with a septum and H2 was added (at a constant flow of 2 atm) over a 15 min period. The reaction
was then stirred for 16 h, before filtration. The filtrate was reduced to 5 mL and then placed in a -35 °C
to grow microcrystals. The filter cake was washed with pentane providing a red solution, which was
removed in vacuo to provide the majority of the product as a red powder (610 mg, 58%). X-ray quality
crystals were grown from a pentane:THF (15:1) solution at -30 °C. 1H NMR (THF-d8):  7.69 (br,
CCH2CH2, 8H), 7.33 (m, CCH2, 8H), 1.54 (CH3, 24H), -15.72 (quintet, CoH, 2JPH = 24 Hz, 1H). 13C{1H}
NMR (THF-d8): 153.4 (PC), 128.8 (CCH2CH2), 127.2 (CCH2), 26.1 (CH3). 31P{1H} NMR (THF-d8): 
39.4 (br). Anal. Calcd for C20H33CoP4: %C, 52.64; %H, 7.29. Found %C, 52.75; %H, 7.14.
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A similar procedure was used to synthesize Co(BPE5)2H and Co(depe)2H; see the SI for full
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[Rh(dmpbz)2](CF3SO3). To a 20 mL vial was added [Rh(COD)2](CF3SO3) (375 mg, 0.801
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mmol) and THF (5 mL). To this solution was slowly added dmpbz (325 mg, 1.64 mmol). The solution
was stirred at room temperature. During this time the solution changes from red to orange and eventually
a yellow precipitate forms. The yellow precipitate is filtered and then washed with 5 mL ether, 5 mL
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pentane and 5 mL THF to remove free COD. Yield: 489 mg, 94%. H NMR (THF-d8):  8.10 (br,
CCH2CH2, 8H), 7.69 (m, CCH2, 8H), 1.89 (CH3, 24H). 13C{1H} NMR (THF-d8): 145.4 (PC), 132.5
(CCH2CH2), 130.2 (CCH2), 18.1 (CH3). 31P{1H} NMR (THF-d8):  33.9 (d, JRhP = 121 Hz). 19F{1H}
NMR (THF-d8):  -77. Anal. Calcd for C21H32RhF3O3SP4: %C, 38.90; %H, 4.97. Found %C, 38.77; %H,
4.81.
A similar procedure was used to synthesize [Rh(depe)2](CF3SO3) and [Rh(dcpe)2](CF3SO3); see the
SI for full details.
Rh(dppe)2H.61 To a solution of [Rh(dppe)2](CF3SO3) (180 mg, 0.172 mmol in 5 mL CH3CN)
was added KOtBu (27 mg, 0.240 mmol in 5 mL CH3CN). This solution was filtered and then H2 was
bubbled through the solution for 10 min. Small orange crystals formed within 2 h that were collected,
1
washed with CH3CN, and then dried. Yield: 96 mg (62%). H NMR (THF-d8):  7.32 (br s, PCCH, 16H),
7.06 (t, J = 10 Hz, PCCHCH, 8H), 6.97 (t, J = 10 Hz, PCCHCHCH, 16H), 2.12 (br s, CH2CH2, 8H), -
10.55 (m, RhH, 1H). 13C{1H} NMR (THF-d8): 143.1 (PC), 132.8, 128.0, 127.8, 32.2 (PCH2) . 31P{1H}
NMR (THF-d8):  56.7 (d, JRhP = 143 Hz). Anal. Calcd for C52H49RhP4: %C, 69.34; %H, 5.48. Found
%C, 69.17; %H, 5.36.
Catalytic Procedures: Acquisition of NMR Spectra: Operators of high-pressure equipment such as
that required for these experiments should take proper precautions to minimize the risk of personal injury.
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1H NMR spectra were collected at 500 MHz on a Varian with a 13 degree pulse, 2 sec acquisition
time and a 1 sec recycle time, total acquisition time for 1 transient was 3 sec. Either 8 or 16 transients
were collected each cycle for a total NMR acquisition time of 24 or 48 sec. Spectra were referenced to
internal residual protonated solvent. These parameters were found to yield highly reproducible
quantitative results under the reaction conditions listed through testing with known concentrations of
formate-containing solutions. At high catalyst concentration the sealed standard containing cobalt(II) and
HDO was integrated against the cobalt-hydride resonance before adding gases. At low catalyst
concentration the hydride resonance was not observable and an external calibration of the sealed HDO
standard was used. As an additional internal standard the area of the proton NMR resonance of the acidic
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proton on the Verkade’s base (which is split into two resonances by scalar coupling from P) was also
used. During kinetic runs no shimming was used before spectral collection and the sample was run without
spinning to minimize the time between cycles. The time quoted for each cycle was from the start of the
first transient. Turnover frequencies were calculated from the slope of the linear portion of the kinetic
curves such as those displayed in the SI. The TOF’s quoted in Table 1 were the average of two to three
experiments run under identical conditions. 31P{1H} NMR spectra were collected at time zero and after
catalysis was complete and for select kinetic runs. Spectra were referenced to external phosphoric acid.
Representative preparation in a J Young glass tube: A glass capillary containing the standard was
added to the tube. Vkd base was dissolved in 200 µL THF-d8 and then added to the tube, followed by 50
µL of the catalyst at the desired dilution. The volume was then made to a total volume of 500 µL with
THF-d8. The tube was sealed and transferred to the gas manifold line. The tube was exposed to a static
vacuum to eliminate N2 from the tube, then the CO2/H2 gas mix was delivered and mixed for 30 sec on a
vortex mixer before collecting an NMR spectrum. The tube was then connected back to the gas manifold
line, fresh gas was added, the sample was mixed for 30 s, and another NMR spectrum was collected. This
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cycle was repeated until the reaction was complete. The length of time elapsed between each refill is
approximately two minutes.
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Representative preparation in a PEEK tube: A glass capillary containing the standard was added
to the tube. The base was dissolved in 150 µL THF-d8 and then added to the tube, followed by 50 µL of
the catalyst at the desired dilution. The volume was then made to a total volume of between 330 and 390
µL with THF-d8. The tube was then sealed and transferred to the ISCO pump apparatus. The tube was
exposed to a static vacuum to eliminate N2 from the tube. Then the CO2/H2 gas mix was delivered at a
constant pressure by an ISCO pump at the desired pressure. The reaction was placed on a vortex mixer
to allow for gas mixing between collection of the NMR spectra. NMR spectra were collected between
two and five minutes apart.
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X-ray Structural Determination. Crystals of [Co(depe)2(CH3CN)](BF4)2 and trans-
[Co(dmpbz)2(CH3CN)2](BF4)2 were grown by ether diffusion into acetonitrile, and crystals of
Co(dmpbz)2H were grown from a 15:1 solution of pentane:THF. Crystals selected for diffraction studies
were immersed in Paratone-N oil, placed on a Nylon loop, and transferred to a precooled cold stream of
N2. A Bruker KAPPA APEX II CCD diffractometer with 0.71073 Å Mo K radiation was used for
diffraction studies. The space groups were determined on the basis of systematic absences and intensity
statistics. The structures were solved by direct methods and refined by full-matrix least squares on F2. All
nonhydrogen atoms were refined anisotropically unless otherwise stated. Hydrogen atoms were placed at
idealized positions and refined using the riding model unless otherwise stated. Data collection and cell
refinement were performed using Bruker APEX2 software. Data reductions and absorption corrections
were performed using Bruker’s SAINT and SADABS programs, respectively.91-92 Structural solutions and
refinements were completed using SHELXS-97 and SHELXL-97,93 respectively, using the OLEX2
software package94 as a front-end.
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Notes on the structure of Co(dmpbz)2H: The complex, which does not possess true two-fold
symmetry, is disordered about a two-fold rotation axis. As such, although there are two
crystallographically distinct dmpbz ligands, they are forced to almost totally overlap by the two-fold axis.
AFIX 66 constraints were placed on the phenyl rings and the SIMU restraint was globally applied.
Although all non-H atoms were refined anisotropically, it was necessary to constrain the thermal ellipsoids
of two sets of methyl groups to be identical using the EADP command. The two sets of ipso carbons of
the dmpbz ligands were also constrained to be identical by the EADP command. The hydride ligand could
not be identified in the Fourier difference map; based on the geometry around Co, which itself appears to
have pseudo two-fold symmetry (distinct from the crystallographic two-fold), the hydride ligand may be
disordered across two positions for each Co environment (thus disordered across four positions in total).
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Notes
The authors declare no competing financial interests.
Supporting Information (SI) Available: Additional synthetic procedures, multinuclear NMR data,
representative kinetic plots, cyclic voltammograms, and thermochemical data (PDF); X-ray
crystallographic data (cif)
Acknowledgements
The research by M.S.J., A.M.A., E.S.W., and J.C.L. was supported by the U.S. Department of
Energy, Office of Science, Office of Basic Energy Sciences, Division of Chemical Sciences, Geosciences,
and Biosciences. The research by E.B.H., M.L.H., and M.T.M. (X-ray crystallography, synthesis) was
supported as part of the Center for Molecular Electrocatalysis, an Energy Frontier Research Center funded
by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences. The authors thank
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Dr. Samantha A. Burgess for assistance in collecting cyclic voltammetry data. Pacific Northwest National
Laboratory is operated by Battelle for the U.S. Department of Energy.
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