Thermodynamic Hydricities of Biomimetic Organic Hydride Donors

Article abstract of DOI:10.1021/jacs.7b13526

Thermodynamic hydricities (ΔG_H^-) in acetonitrile and dimethyl sulfoxide have been calculated and experimentally measured for several metal-free hydride donors: NADH analogs (BNAH, CN-BNAH, Me-MNAH, HEH), methylene tetrahydromethanopterin analogs (BIMH, CAFH), acridine derivatives (Ph-AcrH, Me₂N-AcrH, T-AcrH, 4OH, 2OH, 3NH), and a triarylmethane derivative (6OH). The calculated hydricity values, obtained using density functional theory, showed a reasonably good match (within 3 kcal/mol) with the experimental values, obtained using ;potential pK_a and ;hydride-transfer methods. The hydride donor abilities of model compounds were in the 48.7-85.8 kcal/mol (acetonitrile) and 46.9-84.1 kcal/mol (DMSO) range, making them comparable to previously studied first-row transition metal hydride complexes. To evaluate the relevance of entropic contribution to the overall hydricity, Gibbs free energy differences (ΔG_H^-) obtained in this work were compared with the enthalpy (ΔH_H^-) values obtained by others. The results indicate that, even though ΔH_H^- values exhibit the same trends as ΔG_H^-, the differences between room-temperature ΔG_H^- and ΔH_H^- values range from 3 to 9 kcal/mol. This study also reports a new metal-free hydride donor, namely, an acridine-based compound 3NH, whose hydricity exceeds that of NaBH₄. Collectively, this work gives a perspective of use metal-free hydride catalysts in fuel-forming and other reduction processes.

Full text of DOI:10.1021/jacs.7b13526

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Thermodynamic Hydricities of Biomimetic Organic Hydride Donors
Stefan Ilic, Usha Pandey Kadel, Yasemin Basdogan, John A. Keith, and Ksenija D. Glusac
J. Am. Chem. Soc., Just Accepted Manuscript • DOI: 10.1021/jacs.7b13526 • Publication Date (Web): 16 Mar 2018
Downloaded from http://pubs.acs.org on March 17, 2018
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Thermodynamic Hydricities of Biomimetic Organic
Hydride Donors
Stefan Ilic‡, Usha Pandey Kadel§, Yasemin Basdogan♭, John A. Keith♭, Ksenija D.
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Glusac‡*
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‡Department of Chemistry, University of Illinois at Chicago, Chicago, IL 60607, United States
‡Chemical Sciences and Engineering Division, Argonne National Laboratory, Lemont, IL 60439, United
States
§Department of Chemistry, Center for Photochemical Sciences, Bowling Green State
University, Bowling Green, OH 43403, United States
♭Department of Chemical and Petroleum Engineering, University of Pittsburgh, Pittsburgh, PA 15260,
United States
ABSTRACT: Thermodynamic hydricities (ꢀGH–) in acetonitrile and dimethyl sulfoxide have been calculated
and experimentally measured for several metalꢀfree hydride donors: NADH analogs (BNAH, CNꢀBNAH, Meꢀ
MNAH, HEH), methylene tetrahydromethanopterin analogs (BIMH, CAFH), acridine derivatives (PhꢀAcrH,
Me2NꢀAcrH, TꢀAcrH, 4OH, 2OH, 3NH) and a triarylmethane derivative (6OH). The calculated hydricity values,
obtained using density functional theory, showed a reasonably good match (within 3 kcal/mol) with the
experimental values, obtained using “potentialꢀpKa” and “hydrideꢀtransfer” methods. The hydride donor abilities
of model compounds were in the 48.7 – 85.8 kcal/mol (acetonitrile) and 46.9 – 84.1 kcal/mol (DMSO) range,
making them comparable to previously studied firstꢀrow transition metal hydride complexes. To evaluate the
relevance of entropic contribution to the overall hydricity, Gibbs free energy differences (ꢀGH–) obtained in this
work were compared with the enthalpy (ꢀHH–) values obtained by others. The results indicate that, even though
ꢀHHꢀ values exhibit the same trends as ꢀGH–, the differences between roomꢀtemperature ꢀGH– and ꢀHH– values
range from 3 to 9 kcal/mol. This study also reports a new metalꢀfree hydride donor, namely an acridineꢀbased
compound 3NH, whose hydricity exceeds that of NaBH4. Collectively, this work gives a perspective of use
metalꢀfree hydride catalysts in fuelꢀforming and other reduction processes.
INTRODUCTION
Enzymatic redox reactions often rely on organic cofactors, such as reduced nicotine adenine dinucleotide
(NADH) and flavin adenine dinucleotide (FADH2), to perform hydride transfer reactions to substrates such as
carbonyl compounds,1 carbon dioxide,2 flavins (imines),3 and compounds containing activated C=C bonds.4ꢀ5 The
synthetic analogs of these biological “H2ꢀequivalents” have found applications in chemical laboratories,
particularly when asymmetric transformations are desired. In the presence of a chiral coꢀcatalyst, NADHꢀanalogs
serve as regioꢀ and enantioselective reagents for the reduction of imines to amines,6ꢀ8 carbonyl compounds to
alcohols,8ꢀ10 as well as compounds with C=C bonds to the corresponding saturated analogs.8, 11ꢀ12 NADH analogs
have also been applied to the fuel forming reactions. Specifically, a simple pyridinium ion has been investigated
as facilitating the electrocatalytic and photoelectrocatalytic reduction of CO2 to methanol.13ꢀ16 While the
experimental work was not reproduced by others16 and the mechanism of catalysis still remains unclear, the
computational work indicates that the CO2 reduction may occur by a hydride transfer from dihydropyridine, a
close relative of NADH.17ꢀ21 More recently, other nitrogenꢀcontaining organic compounds (imidazoles,22ꢀ23
pyridazine,24 pyridoxine,25 mercaptopteridine,26ꢀ28 dihydrophenanthridine29 and dihydroacridine29) have also been
shown to perform CO2 reduction reactions. Furthermore, NADH analogꢀbased ligands coordinated with redoxꢀ
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active transition metal (Ru,30 Ir31) or other metal ions (Al32) were shown to perform the photocatalytic or
electrocatalytic reduction of CO2 or water.
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9
ꢂ
Thermodynamic hydricity (Δꢀꢁ ) is a useful parameter that is often used to evaluate the hydride donating
ability of a molecule. It is defined as the Gibbs free energy for a hydrideꢀion release from the compound, with
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ꢂ
lower values of Δꢀꢁ indicating better hydride donors:
R − H → Rꢃ + Hꢄ
Δꢀꢁ
(1)
ꢂ
This thermodynamic parameter provides useful information for the potential application of hydrides in
synthetic reductions of C=C, C=N and C=O bonds as well as in fuelꢀforming reductions of protons and CO2. For
this reason, the hydricities of a large number of metalꢀbased hydrides have been extensively studied using
computational and experimental methods for different solvents.33ꢀ39 Although Δꢀꢁ values exhibit strong solvent
effects,40ꢀ42 the majority of reported hydricities were obtained in acetonitrile. These hydricities were found to
ꢂ
ꢂ
vary in a wide 25 – 120 kcal/mol range. Relevant to fuelꢀforming reactions in acetonitrile, hydrides with Δꢀꢁ
below 76 kcal/mol are thermodynamically capable of proton reduction,43ꢀ44 whereas Δꢀꢁ below 44 kcal/mol is
needed for the reduction of CO2 to formate.45 The hydricity studies on metalꢀbased models have shown several
ꢂ
ꢂ
structural factors that influence the Δꢀꢁ values of metal hydrides: (i) the type of the metal used significantly
alters the hydricities of complexes. Within the same row of the periodic table, metals with lower atomic number
give rise to metal complexes with greater hydride donor ability.46 Within the same group, metals in second and
third rows are generally better hydride donors than the firstꢀrow analogues;33,
44, 47ꢀ48
(ii) the structural and
electronic properties of the ligand can also tune the hydricities. For example, the decrease in the ligand bite angle
contributes to the lowering of Δꢀꢁ values.48ꢀ50 Furthermore, the presence of electronꢀdonating substituents on
the ligand decreases the hydricities of metal complexes;44, 51ꢀ53 (iii) the overall charge of the metal complex also
ꢂ
affects the hydricites, with anionic complexes being stronger hydride donors than the corresponding neutral
analogs;54ꢀ56 (iv) solvent drastically affects hydricities, where more polar solvents (such as water) lower Δꢀꢁ
values.37, 40, 57ꢀ61 It is interesting to note that the hydricity values in different solvents do not scale linearly, making
it possible for a certain reaction to be thermodynamically downhill in one solvent, while uphill in another.57
ꢂ
A systematic analysis of over 150 reported hydricity values for metalꢀbased hydride donors has enabled the
discovery of many elegant catalytic systems in which the critical reduction step involves a hydride transfer.36, 62
Despite being widely present in natural systems, metalꢀfree hydrides lack proper thermodynamic evaluation.
Most studies of organic compounds have focused on weaker hydride donors, such as arylꢀsubstituted
carbocations and quinones.63ꢀ65 Among stronger organic donors, thermodynamic hydricities have been reported
56, 63ꢀ64, 66ꢀ68
only for a limited number of model compounds.
The most hydridic donors have found to be radical
anions of organic hydride donors, such as oneꢀelectron reduced dihydroanthracenes and toluenes.68 Even though
these radical anionics showed excellent hydride donating ability, very negative potentials required for their
formation (< ꢀ1.5 V) and lowꢀstability of active hydride species prevent the practical application. As a general
trend, the hydricities of these organic hydride donors can be lowered by increasing the stability of the cation R+
formed upon the hydride transfer, either through aromatic stabilization or by the introduction of electronꢀ
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ꢂ
donating groups. Due to experimental challenges associated with determination of Δꢀꢁ values, the hydricities of
metalꢀfree donors are often reported in terms of two other parameters that can be obtained from relatively simple
experimental measurements: the enthalpy change associated with the hydride release (Δꢅꢁ
)
69ꢀ71 and the hydride
ꢂ
nucleophilicity (N).72ꢀ73 While the reported Δꢅꢁ values allowed a screening of a large number of metalꢀfree
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ꢂ
ꢂ
hydrides, it is not clear whether the entropic contribution (TΔꢆꢁ ) is negligible or persistent for structurally
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different hydride donors. Similarly, the nucleophilicity N is an empirical parameter that provides useful insights
into the kinetics of hydride transfers from model donors, but the correlation between N values and standard
kinetic parameters (such as activation free energy, ΔG≠) is not straightforward.
H H
H H
H
H H
CN
H H
CONH2
EtOOC
COOEt
N
N
N
N
N
CH2Ph
CH2Ph
BNAH
Ph-AcrH
HEH
CN-BNAH
N
Me-MNAH
N
N
H
OO
H
H
H
O
O
O
O
N
N
N
N
4OH
2OH
Me2N-AcrH
R
T-AcrH
N
H
O
O
O
O
O
O
N
N
H
N
N
O
H
H
N
N
R
R
H
N
N
O
R=C8H17
CAFH
BIMH
6OH
3NH
Scheme 1: Structure of organic hydrides: NADH analogs (BNAH, CNꢀBNAH, MeꢀMNAH, HEH), methylene
tetrahydromethanopterin analogs (BIMH and CAFH), acridine (PhꢀAcrH, Me2NꢀAcrH, TꢀAcrH, 4OH, 2OH,
3NH) and triarlymethane (6OH) derivatives.
ꢂ
In this study, we report the calculated and experimental thermodynamic Δꢀꢁ for model organic hydrides
presented in Scheme 1. Some of the model compounds are direct analogs of the enzymatic cofactors NADH1
(model compounds BNAH, CNꢀBNAH, MeꢀMNAH and HEH) and methylene tetrahydromethanopterin H4MPT+
74
(model compounds BIMH and CAFH). Other model compounds are derived from acridine (PhꢀAcrH, Me2Nꢀ
AcrH, TꢀAcrH, 4OH, 2OH, 3NH) and triarlymethane (6OH) frameworks. The calculated values were determined
in two solvents using density functional theory (DFT) and supported by experimental findings obtained using
ꢂ
electrochemical and hydride transfer methods. A comparison of Δꢀꢁ values obtained here with calculated
ꢂ
Δꢅꢁ values indicate a degree of uncertainty associated with the evaluation of hydride strength using enthalpic
ꢂ
ꢂ
Δꢅꢁ . In specific, the entropic contribution (TΔꢆꢁ ) was found to differ significantly for structurally unrelated
hydride donors. The results of our work were also discussed in terms of the structural and electronic factors that
lead to good hydride donor abilities in metalꢀfree models. Importantly, we discovered a new metalꢀfree
compound with strong hydride donating ability: an acridineꢀbased structure 3NH was shown to exceed the
hydride donor abilities of natural and most artificial metalꢀfree hydride donors. Additionally, the cathodic
behavior of the corresponding cation, 3N+ was shown to be reversible, indicating that this compound can be
utilized in catalysis.
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COMPUTATIONAL METHOD
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Hydricity calculations. All calculations related to hydricity were performed using Gaussian 09 package75 with
the resources of the Ohio Supercomputer Center. The geometries of relevant species (R+ and RꢀH) were
optimized at the ωB97XꢀD/6ꢀ311G(d) level of theory with the conductorꢀlike polarizable continuum model
(CPCM) for solvents (acetonitrile and dimethyl sulfoxide).76ꢀ78 The frequency calculations were performed to
confirm the absence of imaginary frequencies. The output files from the frequency calculations provided the
thermal corrections to free energies (ꢇꢀꢈꢋꢉꢉꢊꢌꢊ) for R+ and RꢀH. The structures optimized at the ωB97XꢀD/6ꢀ
311G(d) level were then used to perform a singleꢀpoint energy calculation at the ωB97XꢀD/6ꢀ
311++G(2df,p)/CPCM(ACN or DMSO) level and the electronic energies (ℰꢍꢋꢉꢌ) of R+ and RꢀH were obtained
from these output files.
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The computational method for hydricity calculation was adopted from the previously published study.79 The
hydricity of a model compound RꢀH is defined as the thermodynamic driving force (Δꢀꢁꢄ) for the following
reaction:
R–H → R+ + H– Δꢀꢁꢄ = ꢀꢎꢃ + ꢀꢁꢏꢐ − ꢀꢎꢄꢁ
where individual Gibbs free energies are defined as follow:
ꢀꢎꢃ = ℰꢍꢋꢉꢌ + ꢇꢀꢈꢉꢊꢊ + ꢇꢀꢉ→
ꢋꢉꢌ
∗
(
)
ꢎꢃ
ꢀꢁꢏꢐ = ꢑℰꢍꢒꢓꢋ + ꢇꢀꢈꢒꢉꢓꢊꢋꢊ + ꢇꢀꢔꢋꢏꢉꢌꢐ + ꢇꢀꢉ∗→ꢕꢁꢏꢐ
ꢀꢎꢄꢁ = ℰꢍꢋꢉꢌ + ꢇꢀꢈꢉꢊꢊ + ꢇꢀꢉ→
ꢋꢉꢌ
∗
(
)
ꢎꢄꢁ
where ℰꢍꢋꢉꢌ and ℰꢍꢒꢓꢋ represent electronic energies in solvated and gasꢀphases, ꢇꢀ
and ꢇꢀꢈꢒꢉꢓꢊꢋꢊ are thermal
ꢋꢉꢌ
ꢈꢉꢊꢊ
ꢋꢉꢌ
correction to the Gibbs free energy in solvated and gasꢀphases, ꢇꢀ
is solvation free energy for the hydride
ꢔꢏꢐ
anion and ꢇꢀꢍ →∗ is a standard state correction (the value is ꢇꢀꢉ∗→ = +1.891 kcal/mol for all species that do not
have gaseous standard state).80ꢀ81 Electronic energies and thermal corrections to the Gibbs free energy were
obtained as previously described. To derive ꢀꢁꢏꢐ, the electronic energy (ℰꢒꢓꢋ = −331.14 kcal/mol) and the
ꢍ
ꢒꢓꢋ
ꢈꢉꢊꢊ
thermal correction (ꢇꢀ
= −6.28 kcal/mol) were obtained for gasꢀphase using the ωB97XꢀD/6ꢀ31+G(d,p)
ꢋꢉꢌ
level of theory. The solvation energy ꢇꢀ
for H– was obtained from the thermochemical cycle connecting gas
ꢔꢏꢐ
phase and solution phase oneꢀelectron reduction, as expressed in the following equation:
ꢇꢀꢔꢋꢏꢉꢌꢐ= ꢁꢀ(ꢋꢁꢉ/ꢌꢖꢁꢄ) ꢀ ꢁꢀꢒꢓꢋ + ꢁꢀꢋꢉꢌꢖ
(ꢁ)
(ꢁ/ꢁꢄ)
ꢒꢓꢋ
(ꢁ/ꢁꢄ)
where ꢁꢀ
and ꢁꢀ(ꢋꢁꢉ/ꢌꢖꢁꢄ) represent the Gibbs free energy changes for the one electron reduction of
hydrogen atom in the gas phase and the solution, respectively. ꢇꢀꢒꢓꢋ(ꢅ/ꢅꢄ) is the negative value of the electron
affinity of hydrogen atom (ꢇꢀꢒꢓꢋ(ꢅ/ꢅꢄ) = −17.39 kcal/mol82). ꢁꢀ(ꢋꢁꢉ/ꢌꢖꢁꢄ) was obtained from the experimental
oneꢀelectron potentials ꢗꢘꢍ/ꢘꢄ: using ꢗꢘꢍ/ꢘꢄ = −0.60 V68 and −0.55 V68 for acetonitrile and dimethyl sulfoxide,
ꢁꢀ(ꢋꢁꢉ/ꢌꢖꢁꢄ) values were estimated to be −84.88 kcal/mol and −86.04 kcal/mol for acetonitrile and dimethyl
ꢋꢉꢌꢖ
(ꢁ)
sulfoxide, respectively. ꢁꢀ
represents the solvation energy of hydrogen atom and this value was computed
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using CPCM/6ꢀ311++G(2df,p) and found to be −0.1 kcal/mol in both solvents. Using this procedure, the
computed values for ꢀꢁꢏꢐ were −404.8 kcal/mol and −406.0 kcal/mol for acetonitrile and dimethyl sulfoxide,
respectively.
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.
.
Reduction potential calculations. The first (EoR+/R ) and second (EoR /R−) reduction potentials for our model
compounds were derived from the calculated driving forces (ꢀGR+/R. and ꢀGR /R−), as follows:
.
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Δꢀꢙꢃ/ꢙ. = ℰꢍꢋꢉꢌ + ꢇꢗꢉ→ − ℰꢍꢋꢉꢌ + ꢇꢗꢉ→
∗
∗
(
)
.
(
)
ꢚ
ꢎ
ꢎ
Δꢀꢙ./ꢙꢄ = ℰꢍꢋꢉꢌ + ꢇꢗꢉ→
− ℰꢍꢋꢉꢌ + ꢇꢗꢉ→
∗
∗
(
)
ꢂ
(
)
.
ꢎ
ꢎ
where electronic energies were obtained by performing singleꢀpoint calculations using the B3LYP83ꢀD3BJ84
and ωB97XꢀD377 exchange correlation functionals with the Def2ꢀTZVP85 basis set using the SMD continuum
solvation model86 on fully optimized structures obtained using the BP8687ꢀD3BJ/Def2ꢀSVP85 model chemistry
with ORCA88. The entropic contributions for the reactant and product states were assumed to be similar, which
resulted in their mutual cancellation. The Δꢀ values were then used to calculate the standard reduction potentials
(ꢗ = − ꢛꢝꢞꢜ). The calculated values were referenced to NHE by subtracting 3.92 V80 from computed absolute
potentials.89
In case of second reduction potentials, accuracies of EoR /R− reduction potentials were systematically improved
.
compared to available experiment when a counter ion (K+) was included in both the ꢟ. and ꢟꢄstates, i.e. using
reduction potentials modeled as ꢟ.ꢀK+ and ꢟꢄꢀ K+. It seemed that adding a counter ion stabilizes the anion
relative to the neutral radical, and this resulted in better agreement with experiment due to error cancellation. We
report our best calculated values in Table 1: first reduction potential was calculated using ωB97XꢀD3
calculations, while the second reduction potential is calculated with ωB97XꢀD3 and the counter ion. The
Supporting Information (Table S1) reports all calculated data.
EXPERIMENTAL SECTION
General methods. All chemicals were purchased from commercial suppliers and used without further
1
purification. H NMR spectra were recorded on a Bruker Avance III 500 MHz system. Steadyꢀstate UV/Vis
absorption spectra were recorded on a Varian Cary 50 UVꢀVis spectrophotometer. 1ꢀBenzylꢀ1,4ꢀ
dihydronicotinamide (BNAH) and 10ꢀMethylꢀ9ꢀphenylacridinium perchlorate (PhꢀAcr+) was purchased from
TCI America. Fluorene (FlH), triphenylmethane (Ph3CH), diphenylyldiphenylmethane (DPE) and Superꢀ
Hydride (1M in THF) were purchased from Sigma. NAD+ analogs (6O+,90 4O+,90 2O+,90 3N+,91 TꢀAcr+,90 Me2Nꢀ
Acr+,90 BNA+,92 CNꢀBNA+,92 MeꢀMNA+,93 HE+,94 BIM+95 and CAF+96), NADH analogs (6OH,97 2OH,98 Phꢀ
AcrH,79 CNꢀBNAH,92 BIMH,95 CAFH96), indicator 9ꢀphenylxanthene (XanH99) and nickelꢀcomplex
([Ni(dmpe)2](PF6)2100) were synthesized according to the previously published procedures.
N,N-dimethyl-4-(10-methyl-9,10-dihydroacridin-9-yl)aniline (Me2N-AcrH): Me2NꢀAcr+ (412 mg, 1 mmol)
was dissolved in 5 mL ethanol and cooled in an ice bath. Sodium borohydride (150 mg, 4 mmol, 4 eq) was then
added and color changed to yellow. The reaction mixture was warmed to room temperature and stirred for
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additional 4 hours. Resulting solution was filtered and the precipitate washed with dichloromethane. The filtrate
was extracted with dichloromethane, organic extracts were combined and solvent evaporated. The yellow oil was
dissolved in ethanol and precipitated by addition of water. The yellow precipitate was filtered, washed with cold
1
water and dried under vacuum to yield 115 g (37%) of pure product. Hꢀ NMR (CD3CN, 500 MHz): 7.30ꢀ7.23
(4H, m), 7.05 (2H, d), 7.00ꢀ6.90 (4H, m), 6.59 (2H, d), 5.13 (1H, s), 3.41 (3H, s), 2.77 (6H, s).
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N,N-dimethyl-4-((10-methyl-9,10-dihydroacridin-9-yl)ethynyl)aniline (T-AcrH): TꢀAcr+ (120 mg, 0.27
mmol) was dissolved in 6 mL ethanol and cooled in an ice bath. Sodium borohydride (62 mg, 1.62 mmol, 6 eq)
was added and the reaction mixture was stirred for 2 hours, which resulted in disappearance of deepꢀblue color.
The reaction mixture was then filtered, filtrate disposed and precipitate washed with dichloromethane.
Dichloromethane solution was evaporated yielding in 30 mg of brownish product (33%). 1Hꢀ NMR (CD3CN, 500
MHz): 7.67 (2H, d), 7.38 (2H, d), 7.33 (2H, t), 7.10ꢀ7.05 (4H, m), 6.73 (2H, d), 5.00 (1H, s), 3.46 (3H, s), 2.98
(6H, s).
Cyclic voltammetry. Cyclic voltammetry was performed using a BASi epsilon potentiostat in a VCꢀ2
voltammetry cell (Bioanalytical Systems) using platinum working electrode (1.6 mm diameter, MFꢀ2013,
Bioanalytical Systems), a nonaqueous Ag/Ag+ reference electrode (MFꢀ2062, Bioanalytical Systems) and a
platinum wire (MWꢀ4130, Bioanalytical Systems) as a counter electrode. The spectroscopic grade solvent
DMSO and the electrolyte tetrabutylammonium perchlorate (TBAP) were purchased from Sigma Aldrich and
used as received. Fast scan rate cyclic voltammetry was performed using CHI 600 C potentiostat and platinum
working electrode (CHIꢀ107, CH instruments, 10 µm diameter). In case of TꢀAcr+, the second standard reduction
potential was obtained by oxidation of TꢀAcr¯, which was prepared inꢀsitu from TꢀAcrH and potassiumꢀ
dymsil.101 Electrochemical potentials were converted to NHE by adding 0.548 V to the experimental
potentials.102
pKa determination. The pKa values of the NADH analogs were determined using the indicator anion method
in DMSO.99 Under inert atmosphere, indicators (InH) were added to a solution of potassium dimsyl
(K+CH3SOCH2¯) to generate the indicator anions (In¯). An excess of indicator solution was added to the
K+CH3SOCH2¯ to ensure the complete consumption of the base. The anion concentrations were determined
using recorded absorbance and In¯ extinction coefficients. Then, the colored In¯ solutions were quenched by
addition of small aliquots of organic hydrides solutions in DMSO. The pKa values for the organic hydrides were
determined using the known indicator pKa value and experimentally obtained equilibrium constants of the
reactions between indicator anions and the hydrides. Indicators were chosen to be within two pKa units from the
hydrides and indicator absorbed in visible spectrum where the other species were transparent.99 In case of the
overlapping absorptions of In¯ and deprotonated hydride Rꢀ (Me2NꢀAcr¯), the absorption of Me2NꢀAcr¯ was
subtracted using its extinction coefficient at λmax for indicator In¯, as described in Supporting information. The
pKa values of indicators used in this study are:99 triphenylmethane (Ph3CH, pKa = 30.6) for 4OH,
diphenylyldiphenylmethane (DPE, pKa = 29.4) for Me2NꢀAcrH, 9ꢀphenylxanthene (XanH, pKa = 27.9) for Phꢀ
AcrH and 6OH, fluorene (FlH, pKa = 22.6) for 2OH.
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Hydride transfer studies. The hydricities of selected model compounds were obtained by determining the
equilibrium constant for the hydride transfer to an appropriate acceptor with known hydride affinity. To ensure
that the equilibrium constant can be reached, the reference compounds were selected so that their hydricities are
within 3 kcal/mol of the hydridicities of our model compounds (as estimated from DFT calculations described in
the computational section). The equilibrium concentration ratios of reactants and products were obtained using
1H NMR spectroscopy. The following steps were performed to ensure that the equilibrium was reached: the
progress of the reaction was monitored until the integration of NMR peaks stopped changing. Then, an additional
amount of one of the products was added and the reaction was monitored again until the equilibrium was
reached. Deuterated acetonitrile and DMSO were used as solvents. All reaction mixtures were prepared in the
glove box using dry reagents and airꢀtight NMR tubes.
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Equilibrium of BNAH and 2O+: BNAH (3.8 mg, 0.018 mmol) and 2O+ (8.1 mg, 0.018 mmol) were dissolved
in 0.6 mL of deuterated acetonitrile or DMSO. The equilibrium constant was reached after 14 days in acetonitrile
yielding Keq=9.61 whereas the equilibrium was reached after 19 days in DMSO yielding Keq =1.68. The hydricity
of 2OH in acetonitrile was obtained from Keq and the reported hydricity of BNAH (59 kcal/mol) as reference.56
In case of DMSO, the hydricity of 2OH (58.3 kcal/mol) was calculated by using potentialꢀpKa method and the
obtained value was used as reference to calculate the hydricity of BNAH in DMSO. The 2OH hydricity was 60
kcal/mol in acetonitrile and the hydricity of BNAH was 57.7 kcal/mol in DMSO.
Equilibrium of BNAH and HE+: BNAH (3.8 mg, 0.018 mmol) and HE+ (5.5 mg, 0.018 mmol) were dissolved
in 0.6 mL deuterated acetonitrile or DMSO. In acetonitrile, the equilibrium was reached after 15 days, yielding
Keq= 87.52. In DMSO, the equilibrium was reached after 49 days, yielding Keq= 1.53. The hydricity of HEH in
acetonitrile was obtained from Keq and the reported hydricity of BNAH (59 kcal/mol) as reference.56 The
hydricity of HEH in case of DMSO was obtained from Keq and hydricity of BNAH (57.7 kcal/mol) as reference.
The HEH hydricity was 61.5 kcal/mol in acetonitrile and 58.2 kcal/mol in DMSO.
Equilibrium of [Ni(dmpe)2H]+ and 3N+ or BIM+. [Ni(dmpe)2H]+ was prepared in situ by addition of 1M
SuperꢀHydride (20 ꢁL, 0.020 mmol) to a solution of [Ni(dmpe)2](PF6)2 (16.2 mg, 0.025 mmol) in 0.6 mL
deuterated acetonitrile. To this solution was then added 3N+ (12.7 mg, 0.025 mmol) or BIM+ (8.1 mg, 0.026
mmol). The Keq = 3.33 was obtained after 15 days for 3N+ and Keq= 0.69 was obtained for BIM+ after 11 days.
From these equilibrium constants and reported hydricity of [Ni(dmpe)2H]+ (49.9 kcal/mol),36 we derived
(
)
(
)
∆ꢀꢁꢄ 3ꢠꢅ = 49.2 kcal/mol and ∆ꢀꢁꢄ ꢡꢢꢣꢅ = 50.1 kcal/mol in acetonitrile.
RESULTS AND DISCUSSION
Calculated Hydricities. The hydricities ∆ꢀꢁꢄ of model compounds RꢀH (eq 1) are calculated from the
absolute Gibbs energies of reactant and product states in the appropriate solvation model:
∆ꢀꢁꢄ = ꢀꢎꢃ + ꢀꢔꢏꢐ − ꢀꢎꢄꢁ
(2)
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While the Gibbs energies of solvated R+ and RꢀH species can be calculated reasonably well using the standard
DFT methodology and solvation models, the calculation of absolute Gibbs free energy for the solvated hydride
ion (ꢀꢔꢏꢐ) represents a challenge. One way to overcome this drawback is to calculate the thermodynamic
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parameters for a hydride transfer reaction between RꢀH and a reference hydride acceptor (such as acridinium
cation or pꢀbenzoquinone) whose hydride affinity is known from the experiment.66,
Alternatively, the
103
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ꢀ
ꢔꢏꢐ value can be obtained as a fitting parameter from the experimental hydricities and calculated Gibbs energies
17, 48, 104ꢀ105
ꢀꢎꢃ and ꢀꢎꢄꢁ
.
Unfortunately, ꢀꢔꢏꢐ values derived from these studies are not consistent (for example,
ꢀꢔꢏꢐ values in acetonitrile were reported to be –400.7 kcal/mol,48 –404.7 kcal/mol105 and –412.7 kcal/mol104).
In collaboration with the Krylov group at the University of Southern California, we previously calculated the
hydricity of an acridineꢀbased hydride donor and the obtained value was in excellent agreement with the
experimental hydricity.79 In our approach, the absolute Gibbs energy ꢀꢔꢏꢐ was obtained as the sum of the gasꢀ
ꢒꢓꢋ
ꢔꢏꢐ
ꢋꢉꢌ
phase energy ꢀ
and the solvent contribution ꢇꢀ
:
ꢔꢏꢐ
ꢀꢔꢏꢐ = ꢀꢒꢓꢋ + ꢇꢀꢋꢉꢌ
(3)
ꢔꢏꢐ
ꢔꢏꢐ
ꢒꢓꢋ
ꢔꢏꢐ
The gas phase energy ꢀ
was calculated using DFT, while the solvation energy ꢇꢀꢔꢋꢏꢉꢌꢐ was derived from the
experimental oneꢀelectron reduction potential of hydrogen atom in a solvent of interest65 and the calculated gasꢀ
phase electron affinity of an Hꢀatom (as detailed in the computational section). The ꢀꢔꢏꢐ values obtained in this
way are −404.8 kcal/mol (in ACN) and −406.0 kcal/mol (in DMSO). In the current manuscript, this
computational methodology was used to calculate the hydricities of our model hydrides in two solvents (ACN
and DMSO, Table 1).
Table 1. Calculated standard reduction potentials (vs. NHE) for R+/R. and R./R¯, pKa values for RH and ꢁGHꢀ
for RH in different solvents.
E1 (R+/R.)1
E2 (R./R¯)2
pKa (RH)3
DMSO
ꢁGHꢀ(RH)4
ꢁGHꢀ(RH)4
ACN
85.8
Compound
6OH
0.08
−1.27
−1.48
−1.17
−1.20
−1.42
−1.02
−1.40
−1.50
−1.84
−1.65
−1.63
−1.69
−1.75
30.4
37.6
26.1
25.6
33.1
14.9
27.0
36.0
38.4
45.8
34.1
38.4
30.8
84.1
73.2
72.8
70.3
66.5
64.7
61.1
60.6
58.3
51.3
48.7
48.6
46.9
4OH
−0.61
−0.25
−0.30
−0.69
−0.09
−0.58
−1.04
−0.94
−1.87
−1.24
−1.51
−1.07
75.1
PhAcrH
Me2NꢀAcrH
CNꢀBNAH
TꢀAcrH
2OH
74.9
72.2
68.5
66.6
62.9
HEH
62.5
BNAH
CAFH
60.3
53.2
MeꢀMNAH
BIMH
50.3
50.3
3NH
48.7
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1 Calculated using ωB97XꢀD3/Def2ꢀTZVP and the SMD continuum solvation model.
2 Calculated using ωB97XꢀD3/Def2ꢀTZVP and the SMD continuum solvation model with a K+ ion.
3 Calculated using the equation 4.
4 Calculated using ωB97XꢀD/6ꢀ311++G(2df,p)/CPCM (DMSO or ACN).
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The calculations were also used to estimate the standard reduction potentials and pKa values of relevant species
(Table 1). While E1(R+/R.) values acquired using sole electronic energies showed a reasonable match with
experimental values (Table S1, Supporting Information), the E2(R./R¯) values using standard procedures did not
match experiment well (mean unsigned error = 0.17 V). Since our original calculated E2 values were consistently
too negative compared to experiment, we speculated this was because the calculated free energies of R¯ states
were systematically too unstable regardless of different exchange correlation functionals, continuum solvation
methods, and basis set sizes. As a simple correction and following previous work,106 we added a positively
charges counter ion, K+, into the calculations on the R¯ and R. states, and the resulting E2 values agreed with
available experimental data much better (mean unsigned error = 0.08 V). Adding an analogous counter ion, Cl¯,
to the states needed for the E1 calculations did not improve the agreement of calculated vs. experiment. The
obtained calculated reduction potentials are then used to estimate the pKa values for the model hydride donors
(eq 4). However, the calculated pKa values (Table 1) are not very accurate. The trends of reduction potentials
and pKa values will be discussed in later section when experimental values are introduced.
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ꢚ
ꢂ
ꢧ
/ꢧ
ꢂ) ꢄꢦꢜ
ꢬ /ꢬ
ꢮꢯꢰꢱꢲꢳꢴ
ꢭ
ꢬ
ꢭ
ꢦꢜ ꢄꢨꢩ.ꢍꢪ ( ꢫ
ꢧꢂ
ꢃ ꢫ
.
.
ꢚ
/ꢬ
ꢵ.ꢩꢪꢶ
ꢤꢥꢓꢎꢁ
=
(4)
Experimental Hydricities: Two experimental approaches were used to determine the hydricities of model
compounds: the “potentialꢀpKa” and “hydride transfer” methods.36 The “potentialꢀpKa method” uses the relevant
standard reduction potentials and pKa values to determine the hydricity of a model compound, as follows:
ꢇꢀꢫꢸꢵ = −ꢹꢗꢍ .= −23.06 ꢗꢍ
ꢎ /ꢎ
ꢟꢃ + ꢷꢄ → ꢟ.
ꢟ. + ꢷꢄ → ꢟꢄ
ꢟꢄ + ꢅꢃ → ꢟꢅ
ꢚ
ꢚ
.
ꢎ
/ꢎ
ꢇꢀꢫꢸꢨ = −ꢹꢗꢎꢍ /ꢎ
.
ꢂ
= −23.06 ꢗꢎꢍ /ꢎ
.
ꢂ
ꢇꢀꢺꢸ = ꢟꢻꢼꢽꢥꢾ = ¯1.364 pꢥꢓꢎꢁ
ꢚ
ꢂ
ꢚ
ꢂ
/ꢁ
/ꢁ
ꢇꢀꢿꢁꣀꣁꣂ = 69.9 ꣃ꣄ꢾꢼ/ꣅ꣆ꢼ; ꢇꢀ꣇ꢁ꣈꣉ = 54.7 ꣃ꣄ꢾꢼ/ꣅ꣆ꢼ
ꢅꢃ + 2ꢷꢄ → ꢅꢄ
ꢚ
ꢂ
/ꢁ
ꢇꢀꢁꢄ = ꢇꢀꢫꢸꢵ + ꢇꢀꢫꢸꢨ + ꢇꢀꢺꢸ + ꢇꢀꢁ
RH → ꢟꢃ + ꢅꢄ
ꢿꣀꣁꣂ/꣇꣈꣉
where ꢇꢀꢫꢸꢵ and ꢇꢀꢫꢸꢨ represent Gibbs free energy changes for first and second electron reduction of NAD+
analogs, calculated using their reduction potentials ( ꢗꢎꢍ .and ꢗꢍ
.
ꢂ
); ꢇꢀꢺꢸ represent a Gibbs free energy
ꢚ
ꢎ /ꢎ
/ꢎ
ꢚ
ꢂ
/ꢁ
change for protonation of R−, calculated using the pKa values of NADH analogs (pꢥꢓꢎꢁ); ꢇꢀꢿꢁꣀꣁꣂ and
ꢚ
ꢇꢀ꣇ꢁ꣈꣉ ꢂrepresent Gibbs free energy changes for two electron reduction of the proton in dimethyl sulfoxide and
/ꢁ
acetonitrile, respectively, using the derived potentials for proton reduction in these solvents;68 ꢇꢀꢁꢄrepresents
hydricity of the studied NADH analog.
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The “hydride transfer” method involves the determination of the equilibrium constant (KHT) for a hydride
transfer from a model hydride RꢀH and a reference hydride acceptor (A+) with known affinity, as follows:
RH + ꣊ꢃ ⇌ ꢟꢃ + ꣊ꢅ
ꢇꢀꢁꢸ = −ꢟꢻꢼꢽꢥꢁꢸ
ꢇꢀꢁꢄ(꣊ꢅ)
AH → ꣊ꢃ + ꢅꢄ
9
RH → ꢟꢃ + ꢅꢄ
ꢇꢀꢁꢄ ꢟꢅ = ꢇꢀꢁꢄ ꣊ꢅ + ꢇꢀꢁꢸ
(
)
(
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where ꢇꢀꢁꢸ represents a Gibbs free energy change for the hydride transfer between examined and referent
hydrides, calculated from experimentally obtained equilibrium constant for the hydride transfer, ꢥꢁꢸ
;
(
)
ꢇꢀꢁꢄ(꣊ꢅ) and ꢇꢀꢁꢄ ꢟꢅ represent hydricities of the reference and examined hydrides.
Both of these approaches have limitations, which necessitated the use of “potentialꢀpKa method” for those
model compounds that exhibited measurable reduction potentials and pKa values in DMSO. On the other hand,
the “hydride transfer” method was used for the model compounds that reached the equilibrium point when
reacted with the reference hydride acceptor.
a)
Potential-pKa method. While this experimental approach is relatively simple, it is limited to the model
compounds whose reduction potentials ꢗꢎꢍ . and ꢗꢍ
are within the electrochemical window of the electrolyte
–
ꢚ
.
/ꢎ
ꢎ /ꢎ
solution (~ –1.9 V vs. NHE for DMSO using TBAP as electrolyte and platinum working electrode). Furthermore,
the pKa values of hydride donors RꢀH can be experimentally determined only if RꢀH is more acidic than the
solvent (for DMSO, pKa=35, which limits the pKa determination for compounds with pKa values lower than
32).107 The calculated pKa values in Table 1 indicate that the acidity of CNꢀBNAH, MeꢀMNAH, HEH, 4OH,
BNAH, BIMH, CAFH are higher than that of the DMSOꢀlimit, indicating that their hydricities are not likely to
be determined using the “potentialꢀpKa” method. However, this argument should be taken loosely due to the low
accuracy of calculated pKa values. Similarly, the reduction potentials ꢗꢎꢍ . and ꢗꢍ
.
ꢂ
need to be less negative
ꢚ
ꢎ /ꢎ
/ꢎ
than the cathodic electrochemical window of the solvent (~ ꢀ1.9 V vs. NHE for DMSO using TBAP as
electrolyte and platinum working electrode).
Standard reduction potentials were obtained using cyclic voltammetry (Figure 1). At low sweep rates, the first
reduction step of 6O+ and most acridineꢀbased models (PhꢀAcr+, Me2NꢀAcr+, 4O+, 2O+ and 3N+) exhibited
reversible electrochemical behavior, indicating a good chemical stability of the corresponding radicals. On the
other hand, the first reduction potentials of pyridinium (CNꢀBNA+, HE+, BNA+ and MeꢀMNA+) and imidazolium
(BIM+ and CAF+) models appear at more negative potentials and are chemically irreversible, possibly due to
radical dimerization.108ꢀ111 The stability of pyridineꢀbased radicals can be increased by the introduction of
substituents in the 4ꢀposition.110 The lower reactivity of acridineꢀbased radicals over the pyridineꢀbased structures
is likely due to higher delocalization of the unpaired spin in the acridineꢀbased radicals.109 The reduction of Tꢀ
Acr+ becomes chemically reversible only at high scan rates (2 kV/s, Figure 1), despite the fact that the compound
is acridineꢀbased. The origin of the chemical irreversibility has not been explored further, but it is interesting to
note that the reduction of the neutral acridineꢀbased analog to form the radical anion is chemically reversible
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even at 100 mV/s.112 The chemical instability of imidazolium radicals has been previously attributed to either
their dimerization111 or to the loss of Hꢀatom and formation of carbene analogs.111, 113ꢀ115 While the oneꢀelectron
reduced CAF+ can form carbene analogs by a loss of a hydrogen atom from the carbon located between two Nꢀ
centers, it is not clear whether BIM+ can undergo similar chemistry by a loss of a phenyl radical. The chemical
reversibility for the oneꢀelectron reduction of pyridinium and imidazolium models could not be achieved (scan
rates up to 10 kV/s were investigated), which prevented us from obtaining the standard reduction potentials for
these processes.
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Figure 1. Cyclic voltammograms of model compound (cations) in the cathodic range: Pt working electrode, Pt
counter electrode, and nonaqueous Ag/AgNO3 reference electrode. Sweep rate, 0.1 V/s (3N+, CNꢀBNA+, HE+,
BNA+, MeꢀMBNA+, BIM+ and CAF+), 25 V/s (6O+, TꢀAcr–, Me2NꢀAcr+, 2O+), 2 kV/s (TꢀAcr+), 100 V/s (Phꢀ
Acr+, 4O+); electrolyte: 0.1 M TBAP in DMSO. The second reduction peak of TꢀAcr+ was obtained from the
oxidation of TꢀAcr– which was formed in situ by the deprotonation of TꢀAcrH in the presence of dimsyl anion
(pKa= 35).107
The second reduction peak was obtained only for NAD+ analogs whose first reduction peaks were reversible
(6O+, PhꢀAcr+, Me2NꢀAcr+, 4O+, 2O+ and 3N+). At low scan rates (100 mV/s), second reduction peaks were
irreversible, likely due to the protonation of the generated anion to form NADH analogs.116 This assignment is
consistent with this assignment is the fact that the reactivity of NAD– anions (reversibility of the second
reduction peak) correlates well with pKa values of the corresponding RꢀH analogs. For example, the second
reduction peak of 6O+ becomes reversible at relatively low scan rates (25 V/s), which is consistent with relatively
low basicity of 6O– anion (the pKa of 6OH is 26.9, see text below). On the other hand, the reversibility for 4O+
requires the scan rates of 100 V/s and the pKa of 4OH is 30.3 (Table 2). In the case of TꢀAcr+, the standard
reduction potential for this process was obtained by electrochemical oxidation of TꢀAcr¯ anion, formed by the
deprotonation of TꢀAcrH (Figure 1). Similar approach was attempted on the compounds that lacked second
reduction peaks, but the experiment was not successful because dimsylꢀbase was not sufficiently strong base to
deprotonate hydrides.
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Figure 2. Indicators and their pKa values (left) used to determine the pKa values of NADH analogs (right):
triphenylmethane (Ph3CH), diphenylyldiphenylmethane (DPE), 9ꢀphenylxanthene (XanH) and fluorene (FlH).107
The experimental pKa values for selected hydrides in DMSO were obtained using the spectrophotometric
method developed by Bordwell.99 The indicators with known pKa values were deprotonated using the dimsyl
anion and then reacted with model NADH analogs. The equilibrium constant for the proton transfer between the
NADH analog and indicator anion was determined by monitoring the absorption of indicator anion at a selected
wavelength (see Figure S2 for example). The accuracy of this method is high (0.05 pKa units) if the acidities of
indicator and NADH analog are within 2 pKa units to ensure that the equilibrium is reached.99 For this reason,
more than one indicator was used to determine the pKa values of model NADH analogs (Figure 2). In case of
Me2NꢀAcrH, the absorption of deprotonated NADH analog overlapped with the absorption of indicator anion. In
this instance, the absorption contribution due to the deprotonated NADH analog was accounted for, as described
in the Supporting Information. Deprotonation of TꢀAcrH led to unstable products, which prevented us from
determining the pKa value of TꢀAcrH. Also, the same methods could not be applied for the acetonitrile, since the
NADH analogs are even weaker acids in this solvent. The obtained pKa values in DMSO were used, along with
the standard reduction potentials, to determine the hydricities of model NADH analogs and the values are
reported in Table 2.
Table 2. Experimentally obtained ꢗꢎꢍ
and ꢗꢎꢍ /ꢎ (V vs. NHE), ꢤꢥꢓꢎꢁ of NADH model as well as ꢇꢀꢁꢄ
.
ꢂ
ꢚ
.
/ꢎ
(kcal/mol) values derived using “potentialꢀpKa” method in dimethylꢀsulfoxide.
ꢗꢎꢍ
.
ꢂ
ꢍ
ꢤꢥꢓꢎꢁ
26.9
30.4
28.3
29.2
N/A
ꢇꢀꢁꢄ
ꢗꢎ /ꢎ
Compound
ꢚ
.
/ꢎ
6OH
4OH
+0.24
ꢀ0.38
ꢀ0.29
ꢀ0.30
ꢀ0.22
ꢀ0.50
ꢀ0.80
ꢀ1.24
ꢀ1.41
ꢀ1.23
ꢀ1.42
ꢀ1.07
ꢀ1.39
ꢀ1.62
83.5 ± 3
70.2 ± 3
73.5 ± 2
70.1 ± 2
N/A
PhAcrH
Me2NꢀAcrH
TꢀAcrH
2OH
23.4
N/A
58.2 ± 2
N/A
3NH
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Hydride Transfer Method. The hydricities of almost one half of the model NADH analogs presented
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in Scheme 1 could not be obtained using the potentialꢀpKa method (either due to the irreversible reduction
behavior of NAD+ analogs or due to the high pKa values of the corresponding RꢀH). The hydricities of these
model compounds were obtained using the hydrideꢀtransfer method, which was previously used by Dubois to
obtain the hydricities of NADHꢀanalogs (BNAH and CNꢀBNAH) in acetonitrile, using the metalꢀbased hydride
donors as references.56 This study showed that the accurate equilibrium constants can be obtained if the
hydricities of two relevant hydrides differ by less than 3 kcal/mol. In our study, NMR spectroscopy was used to
determine the hydricities of model NADH analogs in two solvents: acetonitrile and dimethyl sulfoxide, and the
results are listed in Table 3. As an example, the NMR spectra for the reaction between BNAH (reference
hydride) and 2O+ in acetonitrile is presented in Figure S3, Supporting Information. Unfortunately, the hydride
transfer method could not be applied to all model compounds, due to the occurrence of unwanted side reactions.
Experimental hydricity values obtained potentialꢀpKa and hydride transfer methods are further validated by
crossꢀreferencing among studied model compounds (see Supporting Information for more details).
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ꢂ
Table 3. Hydricities (Δꢀꢁ , kcal/mol) of selected model NADH analogs obtained using “hydrideꢀtransfer”
method in dimethylsulfoxide and acetonitrile.
ꢂ
ꢂ
Compound
BNAH
HEH
Reference Hydride (Δꢀꢁ )
Δꢀꢁ
2OH (58.2)
BNAH (57.5)
2OH (58.2)
57.5 ± 2
58 ± 2
57.9 ± 2
591 ± 2
60.3 ± 2
61.5 ± 2
61 ± 2
BNAH
2OH
CpRe(NO)(CO)(CHO)1 (55)
BNAH (59) 1
BNAH (59) 1
HEH
CNBNAH (63) 1
CNꢀBNAH
BIMH
BNAH (59) 1
631 ± 2
50.1 ± 2
49.2 ± 2
[Ni(dmpe)2H]+ (49.9)36
[Ni(dmpe)2H]+ (49.9)36
3NH
1Ref. 56
Theory vs. Experiment. The experimental hydricities obtained using “potentialꢀpKa” and “hydride transfer”
methods showed a good match with the calculated values (Figure 3). The calculated values are frequently higher
than the experimental hydricities by up to 3 kcal/mol, likely due in part to the uncertainties associated with the
treatment of the hydride ion solvation in calculations. A similar match between experiment and theory has been
48
reported previously for metalꢀbased hydrides.33,
Our finding that computationallyꢀinexpensive methods
reproduce the experimental hydricities is quite encouraging, particularly in light of technical difficulties
associated with determination of experimental values.
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Figure 3. Comparison between the calculated and experimental hydricities for the NADHꢀanalogs. The value
for PhAcrH was obtained from our ref. 79, while the values for BNAH and CNꢀBNAH were obtained from ref. 56
ꢂ
Entropic contributions. The experimental determination of Δꢀꢁ values can be quite challenging, as shown in
ꢂ
the previous section. On the other hand, the enthalpy for hydride transfer Δꢅꢁ is readily obtained by a simple
calorimetry method, and this method has been used to screen a large number of metalꢀfree hydride donors.
ꢂ
While the enthalpy approach is quite useful, the use of Δꢅꢁ to describe hydride donor ability is justified only
ꢂ
when entropic contributions are negligible. To evaluate the entropic contribution, we compared the Δꢀꢁ and
Δꢅꢁ values obtained by us and others56, 69, 71, 103, 117ꢀ119 in Figure 4. The data clearly show that the entropic
ꢂ
ꢂ
contribution for metalꢀfree hydrides cannot be ignored, with TΔꢆꢁ values ranging from 3.4 to 12.1 kcal/mol.
Thus, the assumption that the entropic contribution can be neglected introduces, on average, a 10% error to the
measurement.
A more accurate approach is to assume that the entropic contribution is constant across a series of structurally
ꢂ
related hydride donors. This assumption was found to be valid for tungsten hydride complexes (TΔꢆꢁ = 3.2
kcal/mol).54 While this assumption drastically reduces the error, the Δꢆꢁ values are not always constant across
ꢂ
ꢂ
series. NADH analogs exhibit TΔꢆꢁ values in a narrow range (4.4 – 5.2 kcal/mol), whereas the acridine
derivatives were found to have values in much wider range (3.4 – 6.5 kcal/mol). Thus, the enthalpic hydricity
ꢂ
Δꢅꢁ can be used to predict trends within the structurally related hydride donors, but it fails to give precise
driving forces for hydride transfer reactions with hydrides of different groups.
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ꢂ
ꢂ
Figure 4. a) Comparison between Δꢅꢁ and Δꢀꢁ for organic and other metalꢀfree hydride donors in
ꢂ
ꢂ
acetonitrile. b) Comparison between ꢻΔꢆꢁ and Δꢀꢁ for organic hydride donors. Hydrides are represented as
follows: calculated values for hydrides studied here (red triangles), calculated values by others (blue circles) and
experimental values obtained elsewhere (green squares).
Structure-Property Relationship. Based on the calculated and experimental data presented here, several
interesting points can be raised regarding the structural factors that control hydricities of organic donors. In
general, the hydride donor ability of RꢀH derivatives improves with an increase in the stabilization of the forming
cation, R+. Consequently, triarylmethane derivative showed the poorest hydride donor ability, where the R+
stabilization is achieved solely through nonaromatic delocalization. The positive charge in R+ can be stabilized
through the inductive effect of electronꢀdonating substituents, as can be observed for the NADHꢀanalogs, where
ꢂ
ꢂ
the replacement of electronꢀwithdrawing amide (BNAH, Δꢀꢁ = 59 kcal/mol) or cyano (CNꢀBNAH, Δꢀꢁ = 63
ꢂ
kcal/mol) groups with the electronꢀdonating methyl group (MeꢀMNAH, Δꢀꢁ = 50 kcal/mol) has a strong effect
ꢂ
on hydricity. In another example, the hydricity of Me2NꢀAcrH (Δꢀꢁ =70.1 kcal/mol) is lower relative to the
ꢂ
derivative without donating group (PhAcrH, Δꢀꢁ = 73.5 kcal/mol).
An interesting class of hydride donors are imidazoles (BIMH and CAFH), which have shown low hydricities
69
ꢂ
(calculated ACN values in the 50ꢀ53 kcal/mol range) in accordance from previously reported Δꢅꢁ
.
The high
hydride donor ability of imidazole derivatives has been previously credited to the specific conformation and an
anomeric effect,120 where neighboring nitrogen centers destabilize the CꢀH bond in RꢀH by donating their lone
pairs to its antibonding orbital. It is interesting to note that imidazoleꢀbased hydride donors have been identified
in a relatively new class of hydrogenases, namely 5,10ꢀmethenyltetrahydromethanopterin (Hmd) hydrogenase.121
In this enzyme, the imidazoleꢀbased hydride donor drives H2 heterolysis reaction, activated by the presence of
Feꢀcontaining cofactor. Since the synthetic procedures for these derivatives are relatively straightforward, it is
likely that new imidazoleꢀbased hydride donors will be reported in the near future.
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Strikingly, the strongest hydride donor among the studied compounds was found to be an acridine derivative,
ꢂ
3NH (calculated Δꢀꢁ = 49 kcal/mol in ACN). This finding is surprising, considering that the increase in the
number of fused aromatic rings leads to the lowering of aromatic stabilization in the corresponding R+ cation. For
example, dihydroꢀpyridines are better hydride donors than dihydroꢀacridines, because of the higher aromatic
character of pyridinium cation.103 In the case of 3N+, the aromatic moiety is stabilized through the positive charge
delocalization via the conjugation with the remainder of the molecule. In addition, a significant structural strain
present in the rigid 3NH further facilitates the loss of a hydride ion, resulting in a highly planar system. The
effects of extended charge delocalization and planarization prevails in 4OH < 2OH < 3NH series where hydricity
in ACN declines from 75 kcal/mol to 49 kcal/mol.
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It is interesting to note that acridineꢀbased hydrides studied here (3NH and 2OH) have comparable hydricity
117
ꢂ
values to other known reducing reagent, such as sodium borohydride (NaBH4, calculated ΔGꢘ = 50 kcal/mol
117
ꢂ
in ACN) and Hantzsch ester (ΔGꢘ = 59 kcal/mol in ACN). Despite being commonly utilized in organic
synthesis, these hydrides are used in stoichiometric rather catalytic amounts, due to challenges associated with
their regeneration.122 The electrochemical recovery of NADHꢀanalogs is complicated by the chemical side
reactions involving oneꢀelectron reduced species (such as radical dimerization).108ꢀ109 Unlike NADHꢀanalogs,
3N+ and 2O+ exhibit electrochemically and chemically reversible reduction potentials (Figure 1), suggesting that
3NH and 2OH can be electrochemically recovered and could possibly serve as renewable hydride donors.
Overall, the hydride donor abilities were improved slightly going from acetonitrile to dimethyl sulfoxide (by ~
2 kcal/mol), which was also predicted by calculations. Such a trend can be explained by the small differences in
their dielectric constants (ε(DMSO) = 47 vs. ε(ACN) = 38), which results in slightly better solvation of formed
charged species (R+ and H¯). A similar solvent trend has been observed for the Niꢀbased hydride,40 whose
ꢂ
ꢂ
hydricites were obtained in acetonitrile (Δꢀꢁ = 57.4 kcal/mol), dimethyl sulfoxide (Δꢀꢁ = 55.5 kcal/mol) and
ꢂ
water (Δꢀꢁ = 30.0 kcal/mol).
Comparison with metal-based analogs. Transition metal hydrides have been identified as important
intermediates in a variety of catalytic fuelꢀforming and other redox reactions in ground and excited state.35, 45, 123ꢀ
127 On the other hand, the metalꢀfree hydrides have not been widely used for this purpose, despite the abundance
of enzymatic catalysis by NADH, FADH2 and other metalꢀfree hydrides.1ꢀ5, 121 Thus, it is interesting to compare
the thermodynamic hydricities of our model compounds with those reported for metalꢀbased hydrides. In general,
the hydricities of metalꢀbased hydrides span a wide range of values (reported acetonitrile hydricities are in the
26ꢀ120 kcal/mol range34, 36). The metalꢀfree hydrides of Scheme 1 exhibit values that are somewhat higher, with
the calculated values in the 49ꢀ86 kcal/mol range, and seem to have hydricities similar to the first raw transition
metal hydrides, such as Co, Ni and Feꢀbased compounds (acetonitrile hydricities are in the 32ꢀ73 kcal/mol
range36). Thus, both types of compounds are sufficiently strong hydride donors for the relevant fuel forming
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reactions. For example, the hydride affinity of protons in acetonitrile is ~ –76.6 kcal/mol,43 indicating that most
metalꢀfree hydride donors in Scheme 1 are thermodynamically capable of driving the hydrogen evolution
reaction.
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Why are then metalꢀfree hydride donors not used in fuelꢀforming reactions as often as their metalꢀbased
analogs? One possible explanation might be related to the differences in the activation barriers associated with
the relevant hydride transfer processes (our future studies will investigate the kinetic effect in more detail).
Another reason for lower use of metalꢀfree hydride catalysts might be related to the closure of catalytic cycle,
which involves the twoꢀelectron, protonꢀcoupled reduction of R+ to recover the active RꢀH hydride form. To
exemplify this point, Figure 5 presents an energy diagram for two hydride donors of similar hydricities: a metalꢀ
based [Ni(PPh2NPh)2H]+ complex, whose hydricity in acetonitrile is Δꢀꢁ = 59.3 kcal/mol,128 and 2OH, whose
ꢂ
hydricity is Δꢀꢁ = 60.3 kcal/mol. The first reduction potentials of the corresponding precursors, E(M2+/M+) and
ꢂ
E(R+/R.), are relatively similar and affordably small (around – 0.5 V vs. NHE). Previous studies of metalꢀbased
compounds have shown that E(M2+/M+) shifts to more negative values as the hydride donor ability of the
corresponding donor increases.36, 52, 129ꢀ131 Our metalꢀfree analogs scale in the same way, as exemplified by the
similarities in the first reduction peaks for M2+/M+ and R+/R.. However, a striking difference was observed in the
values for the second reduction potentials, E(M+/M0) and E(R./R–): while the metalꢀbased compound undergoes
the second reduction at a relatively low potential (–0.5 V vs. NHE), the metalꢀfree model compound requires a
significantly more negative potential (–1.4 V vs. NHE) to inject the second electron.
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Figure 5. The comparison of energy diagram profiles for the regeneration of active hydride forms of metalꢀ
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5
based ([Ni(PPh2NPh)2H]+)128 and metalꢀfree (2OH) hydride donors.
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Such a large energy requirement for the second reduction step prohibits the application of metalꢀfree hydride
donors in catalysis. Several approaches can be used to lower the standard reduction potentials in metalꢀfree
systems. One involves the coupling of the first electron transfer step with a proton transfer to generate RH.+,
which will then be reduced at a less negative potential. Such protonꢀcoupled reduction has been selected by
nature as a way to regenerate NADH. Specifically, NAD+ reduction is mediated by FADH2, which is formed
from FAD+ through two protonꢀcoupled reductions.132 Another interesting approach towards the lowering of
reduction potentials has recently been reported by Berben and coworkers, who utilized the coordination with Al
ions to lower the reduction potentials of imineꢀbased ligands.32 Finally, model compounds that exhibit small or
even inverted differences in reduction potentials can be utilized to facilitate the regeneration of metalꢀfree
hydrides.133 For example, it was shown for some organic compounds that the structural changes that accompany
the first electron reduction can result in the lowering of their LUMO orbital energies and associated second
reduction potentials.134 We observe similar effects in the case of BIM+, which exhibits the smallest energy
difference between the calculated E(R+/R.) and E(R./R–) potentials (Table 1), likely brought about by the rotation
of the phenyl ring upon one electron reduction (Figure S4, Supporting Information).
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CONCLUSIONS
Despite their importance in enzymatic redox reactions, metalꢀfree hydride donors are not widely used as
catalysts for fuelꢀforming and other reduction processes. To explore their applicability in catalysis, we
investigated hydricities of several model compounds that are direct analogs of the enzymatic cofactors NADH
and methylene tetrahydromethanopterin, as well as the synthetic hydrides derived from acridine and
triarylmethane frameworks. The hydride donor ability for the model compounds reported here were found to be
similar to metalꢀbased hydrides and dependent on structural motifs, such as size of conjugated molecular
framework, aromaticity and presence of electronꢀdonating groups. Unlike the metalꢀbased equivalents, the metalꢀ
free hydrides exhibited high values for their second reduction potentials, prohibiting the catalyst recovery. Future
design of these hydride donors will be focused on the lowering the reduction potentials and enabling the facile
hydrideꢀform regeneration.
SUPPORTING INFORMATION
The Supporting Information is available free of charge on the ACS Publications website at
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1H NMR spectra, coordinates for optimized structures, calculated reduction potentials, pKaꢀ
determination experiment and hydrideꢀtransfer method, enthalpic contributions.
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AUTHOR INFORMATION
Corresponding Author
* glusac@uic.edu
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ORCID
Ksenija D. Glusac: 0000ꢀ0002ꢀ2734ꢀ057X
John A. Keith: 0000ꢀ0002ꢀ6583ꢀ6322
Stefan Ilic: 0000ꢀ0002ꢀ6305ꢀ4001
ACKNOWLEDGEMENTS
K.D.G. thanks ACS PRF (54436ꢀND4) for financial support and Ohio Supercomputer Center (PCS0201ꢀ5) for
computational support. J.A.K. acknowledges support from the R.K. Mellon Foundation, ACS PRF (55595ꢀDNI4)
and the National Science Foundation (CBETꢀ1653392).
Reference
1.
2.
Hummel, W., Trends Biotechnol. 1999, 17 (12), 487ꢀ492.
Reda, T.; Plugge, C. M.; Abram, N. J.; Hirst, J., Proc. Natl. Acad. Sci. U.S.A.
2008, 105 (31), 10654ꢀ10658.
3.
Schrittwieser, J. H.; Velikogne, S.; Kroutil, W., Adv. Synth. Catal. 2015, 357 (8),
1655ꢀ1685.
4.
Stueckler, C.; Hall, M.; Ehammer, H.; Pointner, E.; Kroutil, W.; Macheroux, P.;
Faber, K., Org. Lett. 2007, 9 (26), 5409ꢀ5411.
5.
Stuermer, R.; Hauer, B.; Hall, M.; Faber, K., Curr. Opin. Chem. Biol. 2007, 11
(2), 203ꢀ213.
6.
Rueping, M.; Sugiono, E.; Azap, C.; Theissmann, T.; Bolte, M., Org. Lett. 2005,
7 (17), 3781ꢀ3783.
7.
Ouellet, S. G.; Walji, A. M.; Macmillan, D. W., Acc. Chem. Res. 2007, 40 (12),
1327ꢀ1339.
8.
9.
Zheng, C.; You, S.ꢀL., Chem. Soc. Rev. 2012, 41 (6), 2498ꢀ2518.
Rueping, M.; Antonchick, A. P.; Theissmann, T., Angew. Chem. Int. Ed. 2006, 45
(22), 3683ꢀ3686.
10.
11.
Yang, J. W.; List, B., Org. Lett. 2006, 8 (24), 5653ꢀ5655.
Yang, J. W.; Hechavarria Fonseca, M. T.; List, B., Angew. Chem. Int. Ed. 2004,
43 (48), 6660ꢀ6662.
12.
Ouellet, S. G.; Tuttle, J. B.; MacMillan, D. W., J. Am. Chem. Soc. 2005, 127 (1),
32ꢀ33.
13.
Seshadri, G.; Lin, C.; Bocarsly, A. B., J. Electroanal. Chem. 1994, 372 (1), 145ꢀ
150.
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Page 20 of 24
1
2
3
4
5
6
7
8
14.
Barton, E. E.; Rampulla, D. M.; Bocarsly, A. B., J. Am. Chem. Soc. 2008, 130
(20), 6342ꢀ6344.
15.
Polyansky, D.; Cabelli, D.; Muckerman, J. T.; Fujita, E.; Koizumi, T. a.;
Fukushima, T.; Wada, T.; Tanaka, K., Angew. Chem. Int. Ed. 2007, 46 (22), 4169ꢀ4172.
16.
Costentin, C.; Canales, J. C.; Haddou, B.; Savéant, J.ꢀM., J. Am. Chem. Soc. 2013,
9
135 (47), 17671ꢀ17674.
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
17.
Lim, C.ꢀH.; Holder, A. M.; Hynes, J. T.; Musgrave, C. B., J. Am. Chem. Soc.
2014, 136 (45), 16081ꢀ16095.
18.
2015, 6 (24), 5078ꢀ5092.
19.
20.
21.
22.
Lim, C.ꢀH.; Holder, A. M.; Hynes, J. T.; Musgrave, C. B., J. Phys. Chem. Lett.
Keith, J. A.; Carter, E. A., Chem. Sci. 2013, 4 (4), 1490ꢀ1496.
Keith, J. A.; Carter, E. A., J. Phys. Chem. Lett. 2013, 4 (23), 4058ꢀ4063.
Lessio, M.; Senftle, T. P.; Carter, E. A., ACS Energy Lett. 2016, 1 (2), 464ꢀ468.
Liao, K.; Askerka, M.; Zeitler, E. L.; Bocarsly, A. B.; Batista, V. S., Top. Catal.
2015, 58 (1), 23ꢀ29.
23.
Bocarsly, A. B.; Gibson, Q. D.; Morris, A. J.; L’Esperance, R. P.; Detweiler, Z.
M.; Lakkaraju, P. S.; Zeitler, E. L.; Shaw, T. W., ACS Catalysis 2012, 2 (8), 1684ꢀ1692.
24.
Apaydin, D.; Neugebauer, H.; Knör, G.; Sariciftci, N. S., ChemElectroChem 2014, 1 (9),
1543ꢀ1548.
25.
26.
14010.
27.
28.
121 (16), 4158ꢀ4167.
29.
30.
2016, 55 (5), 2076ꢀ2084.
31.
2017 doi: 10.1002/cptc.201700133.
32.
11646.
33.
34.
Portenkirchner, E.; Enengl, C.; Enengl, S.; Hinterberger, G.; Schlager, S.;
Lee, J. H.; Lauw, S. J.; Webster, R. D., Electrochem commun. 2016, 64, 69ꢀ73.
Xiang, D.; Magana, D.; Dyer, R. B., J. Am. Chem. Soc 2014, 136 (40), 14007ꢀ
Saveant, J.ꢀM.; Tard, C. d., J. Am. Chem. Soc. 2016, 138 (3), 1017ꢀ1021.
Lim, C.ꢀH.; Holder, A. M.; Hynes, J. T.; Musgrave, C. B., J. Phys. Chem. B 2017,
Giesbrecht, P. K.; Herbert, D. E., ACS Energy Lett. 2017, 2 (3), 549ꢀ555.
Kobayashi, K.; Ohtsu, H.; Nozaki, K.; Kitagawa, S.; Tanaka, K., Inorg Chem
Sato, S.; Morikawa, T., ChemPhotoChem 2017, Published Online: October 6,
Thompson, E. J.; Berben, L. A., Angew. Chem. Int. Ed. 2015, 54 (40), 11642ꢀ
Qi, X.ꢀJ.; Fu, Y.; Liu, L.; Guo, Q.ꢀX., Organometallics 2007, 26 (17), 4197ꢀ4203.
DuBois, D. L., Inorg Chem 2014, 53 (8), 3935ꢀ3960.
Bullock, R. M.; Appel, A. M.; Helm, M. L., Chem. Comm. 2014, 50 (24), 3125ꢀ
35.
3143.
36.
Wiedner, E. S.; Chambers, M. B.; Pitman, C. L.; Bullock, R. M.; Miller, A. J.;
Appel, A. M., Chem. Rev. 2016, 116 (15), 8655ꢀ8692.
37. Pitman, C. L.; Brereton, K. R.; Miller, A. J., J. Am. Chem. Soc. 2016, 138 (7),
2252ꢀ2260.
38.
Brereton, K. R.; Bellows, S. M.; Fallah, H.; Lopez, A. A.; Adams, R. M.; Miller,
A. J.; Jones, W. D.; Cundari, T. R., J. Phys. Chem. B 2016, 120 (50), 12911ꢀ12919.
39.
Szalda, D. J.; Muckerman, J. T.; Fujita, E., Angew. Chem. Int. Ed. 2015, 54 (47), 14128ꢀ
14132.
Garg, K.; Matsubara, Y.; Ertem, M. Z.; Lewandowska‐Andralojc, A.; Sato, S.;
ACS Paragon Plus Environment

Page 21 of 24
Journal of the American Chemical Society
1
2
3
4
5
6
7
8
9
40.
Tsay, C.; Livesay, B. N.; Ruelas, S.; Yang, J. Y., J. Am. Chem. Soc. 2015, 137
(44), 14114ꢀ14121.
41.
42.
43.
44.
Taheri, A.; Berben, L. A., Inorg Chem 2015, 55 (2), 378–385.
Creutz, C.; Chou, M. H., J. Am. Chem. Soc. 2007, 129 (33), 10108ꢀ10109.
Wayner, D. D.; Parker, V. D., Acc. Chem. Res. 1993, 26 (5), 287ꢀ294.
Curtis, C. J.; Miedaner, A.; Ellis, W. W.; DuBois, D. L., J. Am. Chem. Soc. 2002,
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
124 (9), 1918ꢀ1925.
45.
46.
47.
DuBois, D. L.; Berning, D. E., Appl. Organomet. Chem. 2000, 14 (12), 860ꢀ862.
Labinger, J. A.; Komadina, K. H., J. Organomet. Chem. 1978, 155 (2), C25ꢀC28.
Curtis, C. J.; Miedaner, A.; Raebiger, J. W.; DuBois, D. L., Organometallics
2004, 23 (3), 511ꢀ516.
48.
Nimlos, M. R.; Chang, C. H.; Curtis, C. J.; Miedaner, A.; Pilath, H. M.; DuBois,
D. L., Organometallics 2008, 27 (12), 2715ꢀ2722.
49.
Miedaner, A.; Haltiwanger, R. C.; DuBois, D. L., Inorg Chem 1991, 30 (3), 417ꢀ
427.
50.
Raebiger, J. W.; Miedaner, A.; Curtis, C. J.; Miller, S. M.; Anderson, O. P.;
DuBois, D. L., J. Am. Chem. Soc. 2004, 126 (17), 5502ꢀ5514.
51. Berning, D. E.; Noll, B. C.; DuBois, D. L., J. Am. Chem. Soc. 1999, 121 (49),
11432ꢀ11447.
52.
Berning, D. E.; Miedaner, A.; Curtis, C. J.; Noll, B. C.; Rakowski DuBois, M. C.;
DuBois, D. L., Organometallics 2001, 20 (9), 1832ꢀ1839.
53.
Chem 2010, 49 (8), 3918ꢀ3926.
54. Ellis, W. W.; Ciancanelli, R.; Miller, S. M.; Raebiger, J. W.; Rakowski DuBois,
M.; DuBois, D. L., J. Am. Chem. Soc. 2003, 125 (40), 12230ꢀ12236.
55. Roberts, J. A.; Appel, A. M.; DuBois, D. L.; Bullock, R. M., J. Am. Chem. Soc.
2011, 133 (37), 14604ꢀ14613.
Wilson, A. D.; Miller, A. J.; DuBois, D. L.; Labinger, J. A.; Bercaw, J. E., Inorg
56.
Ellis, W. W.; Raebiger, J. W.; Curtis, C. J.; Bruno, J. W.; DuBois, D. L., J. Am.
Chem. Soc. 2004, 126 (9), 2738ꢀ2743.
57.
58.
Ceballos, B. M.; Tsay, C.; Yang, J. Y., Chem. Comm. 2017, 53 (53), 7405ꢀ7408
Creutz, C.; Chou, M. H., J. Am. Chem. Soc. 2009, 131 (8), 2794ꢀ2795.
59.
Connelly, S. J.; Wiedner, E. S.; Appel, A. M., Dalton Trans. 2015, 44 (13), 5933ꢀ
5938.
60.
Taheri, A.; Thompson, E. J.; Fettinger, J. C.; Berben, L. A., ACS Catalysis 2015,
5 (12), 7140ꢀ7151.
61.
Robinson, S. J. C.; Zall, C. M.; Miller, D. L.; Linehan, J. C.; Appel, A. M., Dalton
Trans. 2016, 45 (24), 10017ꢀ10023.
62.
Waldie, K. M.; Ostericher, A. L.; Reineke, M.; Sasayama, A. F.; Kubiak, C. P.,
ACS Catalysis 2017.
63.
Cheng, J.; Handoo, K. L.; Parker, V. D., J. Am. Chem. Soc. 1993, 115 (7), 2655ꢀ
2660.
64.
Cheng, J.; Handoo, K. L.; Xue, J.; Parker, V. D., J. Org. Chem. 1993, 58 (19),
5050ꢀ5054.
65.
66.
Parker, V., Acta Chem. Scand. 1992, 46 (11), 1133ꢀ1134.
Zhu, X.ꢀQ.; Wang, C.ꢀH.; Liang, H.; Cheng, J.ꢀP., J. Org. Chem. 2007, 72 (3),
945ꢀ956.
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 22 of 24
1
2
3
4
5
6
7
8
67.
4678.
68.
5067ꢀ5072.
69.
Zhang, X.ꢀM.; Bruno, J. W.; Enyinnaya, E., J. Org. Chem. 1998, 63 (14), 4671ꢀ
Handoo, K. L.; Cheng, J. P.; Parker, V. D., J. Am. Chem. Soc. 1993, 115 (12),
Zhu, X.ꢀQ.; Zhang, M.ꢀT.; Yu, A.; Wang, C.ꢀH.; Cheng, J.ꢀP., J. Am. Chem. Soc.
9
2008, 130 (8), 2501ꢀ2516.
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
70.
6529.
71.
Han, X.; Hao, W.; Zhu, X.ꢀQ.; Parker, V. D., J. Org. Chem. 2012, 77 (15), 6520ꢀ
Zhu, X.ꢀQ.; Tan, Y.; Cao, C.ꢀT., J. Phys. Chem. B 2010, 114 (5), 2058ꢀ2075.
Richter, D.; Tan, Y.; Antipova, A.; Zhu, X. Q.; Mayr, H., Chem. Asian J. 2009, 4
72.
(12), 1824ꢀ1829.
73.
74.
Richter, D.; Mayr, H., Angew. Chem. Int. Ed. 2009, 48 (11), 1958ꢀ1961.
Hiromoto, T.; Warkentin, E.; Moll, J.; Ermler, U.; Shima, S., Angew. Chem. Int.
Ed. 2009, 48 (35), 6457ꢀ6460.
75. Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.;
Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; Nakatsuji,
H.; Caricato, M.; Li, X.; Hratchian, H. P.; Izmaylov, A. F.; Bloino, J.; Zheng, G.;
Sonnenberg, J. L.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida,
M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Montgomery Jr., J. A.;
Peralta, J. E.; Ogliaro, F.; Bearpark, M. J.; Heyd, J.; Brothers, E. N.; Kudin, K. N.;
Staroverov, V. N.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A. P.; Burant,
J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, N. J.; Klene, M.; Knox, J.
E.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.;
Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Martin, R. L.;
Morokuma, K.; Zakrzewski, V. G.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.;
Dapprich, S.; Daniels, A. D.; Farkas, Ö.; Foresman, J. B.; Ortiz, J. V.; Cioslowski, J.;
Fox, D. J. Gaussian 09, Gaussian, Inc.: Wallingford, CT, USA, 2009.
76.
77.
6620.
78.
79.
Chai, J.ꢀD.; HeadꢀGordon, M., J. Chem. Phys. 2008, 128 (8), 084106.
Chai, J.ꢀD.; HeadꢀGordon, M., Phys. Chem. Chem. Phys. 2008, 10 (44), 6615ꢀ
Barone, V.; Cossi, M., J. Phys. Chem. A 1998, 102 (11), 1995ꢀ2001.
Yang, X.; Walpita, J.; Zhou, D.; Luk, H. L.; Vyas, S.; Khnayzer, R. S.; Tiwari, S.
C.; Diri, K.; Hadad, C. M.; Castellano, F. N., J. Phys. Chem. B 2013, 117 (49), 15290ꢀ
15296.
80.
422.
81.
Kelly, C. P.; Cramer, C. J.; Truhlar, D. G., J. Phys. Chem. B 2007, 111 (2), 408ꢀ
Muckerman, J. T.; Skone, J. H.; Ning, M.; WasadaꢀTsutsui, Y., Biochim. Biophys.
Acta 2013, 1827 (8), 882ꢀ891.
82.
83.
84.
85.
86.
6378ꢀ6396.
87.
88.
Lykke, K.; Murray, K.; Lineberger, W., Phys. Rev. A 1991, 43 (11), 6104.
Becke, A. D., J. Chem. Phys. 1993, 98 (7), 5648ꢀ5652.
Grimme, S.; Ehrlich, S.; Goerigk, L., J. Comput. Chem. 2011, 32 (7), 1456ꢀ1465.
Weigend, F.; Ahlrichs, R., Phys. Chem. Chem. Phys. 2005, 7 (18), 3297ꢀ3305.
Marenich, A. V.; Cramer, C. J.; Truhlar, D. G., J. Phys. Chem. B 2009, 113 (18),
Perdew, J. P., Phys. Rev. B 1986, 33 (12), 8822.
Neese, F., Wiley Interdiscip Rev Comput Mol Sci 2012, 2 (1), 73ꢀ78.
ACS Paragon Plus Environment

Page 23 of 24
Journal of the American Chemical Society
1
2
3
4
5
6
7
8
89.
90.
Isse, A. A.; Gennaro, A., J. Phys. Chem. B 2010, 114 (23), 7894ꢀ7899.
Ilic, S.; Brown, E. S.; Xie, Y.; Maldonado, S.; Glusac, K. D., J. Phys. Chem. C
2016, 120 (6), 3145ꢀ3155.
91.
92.
W Laursen, B.; C Krebs, F., Chem. Eur. J. 2001, 7 (8), 1773ꢀ1783.
Paul, C. E.; Gargiulo, S.; Opperman, D. J.; Lavandera, I.; GotorꢀFernández, V.;
9
Gotor, V.; Taglieber, A.; Arends, I. W.; Hollmann, F., Org. Lett. 2012, 15 (1), 180ꢀ183.
93. Huang, S.; Wong, J. C.; Leung, A. K.; Chan, Y. M.; Wong, L.; Fernendez, M. R.;
Miller, A. K.; Wu, W., Tetrahedron Lett. 2009, 50 (35), 5018ꢀ5020.
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
94.
95.
96.
Ogata, Y.; Takagi, K.; Tanabe, Y., Perkin Trans. 1979, (8), 1069ꢀ1071.
Lee, C. K.; Lee, I.ꢀS. H., Heterocycles 2009, 78 (2), 425ꢀ433.
Hori, M.; Kataoka, T.; Shimizu, H.; Imai, E.; Matsumoto, Y., Chem. Pharm. Bull.
1985, 33 (9), 3681ꢀ3688.
97.
98.
99.
Carey, F. A.; Tremper, H. S., J. Am. Chem. Soc. 1968, 90, 2578ꢀ2583.
Guin, J.; Besnard, C.; Pattison, P.; Lacour, J., Chem. Sci. 2011, 2 (3), 425ꢀ428.
Matthews, W. S.; Bares, J. E.; Bartmess, J. E.; Bordwell, F.; Cornforth, F. J.;
Drucker, G. E.; Margolin, Z.; McCallum, R. J.; McCollum, G. J.; Vanier, N. R., J. Am.
Chem. Soc. 1975, 97 (24), 7006ꢀ7014.
100. Robinson, S. J. C.; Zall, C. M.; Miller, D. L.; Linehan, J. C.; Appel, A. M., Dalton
Trans. 2016, 45 (24), 10017ꢀ10023.
101. Cheng, J.ꢀP.; Lu, Y.; Zhu, X.; Mu, L., J. Org. Chem. 1998, 63 (18), 6108ꢀ6114.
102. Pavlishchuk, V. V.; Addison, A. W., Inorg. Chim. Acta. 2000, 298 (1), 97ꢀ102.
103. Shi, J.; Huang, X.ꢀY.; Wang, H.ꢀJ.; Fu, Y., J. Chem. Inf. Model. 2011, 52 (1), 63ꢀ
75.
104. Muckerman, J. T.; Achord, P.; Creutz, C.; Polyansky, D. E.; Fujita, E., Proc. Natl.
Acad. Sci. U.S.A. 2012, 109 (39), 15657ꢀ15662.
105. Kovács, G.; Pápai, I., Organometallics 2006, 25 (4), 820ꢀ825.
106. Keith, J. A.; Grice, K. A.; Kubiak, C. P.; Carter, E. A., J. Am. Chem. Soc. 2013,
135 (42), 15823ꢀ15829.
107. Bordwell, F. G., Acc. Chem. Res. 1988, 21 (12), 456ꢀ463.
108. Hapiot, P.; Moiroux, J.; Saveant, J. M., J. Am. Chem. Soc. 1990, 112 (4), 1337ꢀ
1343.
109. Anne, A.; Hapiot, P.; Moiroux, J.; Savéant, J.ꢀM., J. Electroanal. Chem. 1992,
331 (1ꢀ2), 959ꢀ970.
110. Hermolin, J.; Levin, M.; Kosower, E. M., J. Am. Chem. Soc. 1981, 103 (16),
4808ꢀ4813.
111. de Robillard, G.; Devillers, C. H.; Kunz, D.; Cattey, H.; Digard, E.; Andrieu, J.,
Org. Lett. 2013, 15 (17), 4410ꢀ4413.
112. Elangovan, A.; Chiu, H.ꢀH.; Yang, S.ꢀW.; Ho, T.ꢀI., Org. Biomol. Chem. 2004, 2
(21), 3113ꢀ3118.
113. Gorodetsky, B.; Ramnial, T.; Branda, N. R.; Clyburne, J. A., Chem. Comm. 2004,
(17), 1972ꢀ1973.
114. Ogawa, K. A.; Boydston, A. J., Chem. Lett. 2014, 43 (6), 907ꢀ909.
115. Han, X.; Hao, W.; Zhu, X.ꢀQ.; Parker, V. D., The Journal of organic chemistry
2012, 77 (15), 6520ꢀ6529.
116. Koper, N.; Jonker, S.; Verhoeven, J.; Van Dijk, C., Recl. Trav. Chim. Pays-Bas
1985, 104 (11), 296ꢀ302.
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 24 of 24
1
2
3
4
5
6
7
8
117. Heiden, Z. M.; Lathem, A. P., Organometallics 2015, 34 (10), 1818ꢀ1827.
118. Matsubara, Y.; Fujita, E.; Doherty, M. D.; Muckerman, J. T.; Creutz, C., J. Am.
Chem. Soc. 2012, 134 (38), 15743ꢀ15757.
119. Kreevoy, M. M.; Ostovic, D.; Lee, I. S. H.; Binder, D. A.; King, G. W., J. Am.
Chem. Soc. 1988, 110 (2), 524ꢀ530.
9
120. Brunet, P.; Wuest, J. D., J. Org. Chem. 1996, 61 (6), 2020ꢀ2026.
121. Shima, S.; Pilak, O.; Vogt, S.; Schick, M.; Stagni, M. S.; MeyerꢀKlaucke, W.;
Warkentin, E.; Thauer, R. K.; Ermler, U., Science 2008, 321 (5888), 572ꢀ575.
122. McSkimming, A.; Colbran, S. B., Chem. Soc. Rev. 2013, 42 (12), 5439ꢀ5488.
123. Rakowski Dubois, M.; Dubois, D. L., Acc. Chem. Res. 2009, 42 (12), 1974ꢀ1982.
124. Thoi, V. S.; Sun, Y.; Long, J. R.; Chang, C. J., Chem. Soc. Rev. 2013, 42 (6),
2388ꢀ2400.
125. Miller, A. J.; Labinger, J. A.; Bercaw, J. E., Organometallics 2011, 30 (16), 4308ꢀ
4314.
126. Pitman, C.; Finster, O.; Miller, A., Chem. Comm. 2016, 52 (58), 9105ꢀ9108.
127. Barrett, S. M.; Pitman, C. L.; Walden, A. G.; Miller, A. J., J. Am. Chem. Soc.
2014, 136 (42), 14718ꢀ14721.
128. Fraze, K.; Wilson, A. D.; Appel, A. M.; Rakowski DuBois, M.; DuBois, D. L.,
Organometallics 2007, 26 (16), 3918ꢀ3924.
129. Yang, J. Y.; Bullock, R. M.; Shaw, W. J.; Twamley, B.; Fraze, K.; DuBois, M. R.;
DuBois, D. L., J. Am. Chem. Soc. 2009, 131 (16), 5935ꢀ5945.
130. Galan, B. R.; Schöffel, J.; Linehan, J. C.; Seu, C.; Appel, A. M.; Roberts, J. A.;
Helm, M. L.; Kilgore, U. J.; Yang, J. Y.; DuBois, D. L., J. Am. Chem. Soc. 2011, 133
(32), 12767ꢀ12779.
131. Yang, J. Y.; Smith, S. E.; Liu, T.; Dougherty, W. G.; Hoffert, W. A.; Kassel, W.
S.; DuBois, M. R.; DuBois, D. L.; Bullock, R. M., J. Am. Chem. Soc. 2013, 135 (26),
9700ꢀ9712.
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
132. Carrillo, N.; Ceccarelli, E. A., Eur. J. Biochem. 2003, 270 (9), 1900ꢀ1915.
133. Evans, D. H., Chem. Rev. 2008, 108 (7), 2113ꢀ2144.
134. Evans, D. H.; Busch, R. W., J. Am. Chem. Soc. 1982, 104 (19), 5057ꢀ5062.
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