Energy-Efficient Hydrogen Evolution by Fe-S Electrocatalysts: Mechanistic Investigations

Article abstract of DOI:10.1021/acs.inorgchem.8b00543

The intrinsic catalytic property of a Fe-S complex toward H₂ evolution was investigated in a wide range of acids. The title complex exhibited catalytic events at -1.16 and -1.57 V (vs Fc⁺/Fc) in the presence of trifluoromethanesulfonic acid (HOTf) and trifluoroacetic acid (TFA), respectively. The processes corresponded to the single reduction of the Fe-hydride-S-proton and Fe-hydride species, respectively. When anilinium acid was used, the catalysis occurred at -1.16 V, identical with the working potential of the HOTf catalysis, although the employment of anilinium acid was only capable of achieving the Fe-hydride state on the basis of the spectral and calculated results. The thermodynamics and kinetics of individual steps of the catalysis were analyzed by density functional theory (DFT) calculations and electroanalytical simulations. The stepwise CCE or CE (C, chemical; E, electrochemical) mechanism was operative from the HOTf or TFA source, respectively. In contrast, the involvement of anilinium acid most likely initiated a proton-coupled electron transfer (PCET) pathway that avoided the disfavored intermediate after the initial protonation. Via the PCET pathway, the heterogeneous electron transfer rate was increased and the overpotential was decreased by 0.4 V in comparison with the stepwise pathways. The results showed that the PCET-involved catalysis exhibited substantial kinetic and thermodynamic advantages in comparison to the stepwise pathway; thus, an efficient catalytic system for proton reduction was established.

Full text of DOI:10.1021/acs.inorgchem.8b00543

Article
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Cite This: Inorg. Chem. XXXX, XXX, XXX−XXX
Energy-Efficient Hydrogen Evolution by Fe−S Electrocatalysts:
Mechanistic Investigations
Kai-Ti Chu,† Yu-Chiao Liu,*,† Min-Wen Chung,† Agus Riyanto Poerwoprajitno,† Gene-Hsiang Lee,‡
and Ming-Hsi Chiang*,†
†Institute of Chemistry, Academia Sinica, Nankang, Taipei 115, Taiwan
‡Instrumentation Center, National Taiwan University, Taipei 106, Taiwan
S
* Supporting Information
ABSTRACT: The intrinsic catalytic property of a Fe−S complex toward H2 evolution was
investigated in a wide range of acids. The title complex exhibited catalytic events at −1.16 and
−1.57 V (vs Fc+/Fc) in the presence of trifluoromethanesulfonic acid (HOTf) and
trifluoroacetic acid (TFA), respectively. The processes corresponded to the single reduction
of the Fe-hydride-S-proton and Fe-hydride species, respectively. When anilinium acid was used,
the catalysis occurred at −1.16 V, identical with the working potential of the HOTf catalysis,
although the employment of anilinium acid was only capable of achieving the Fe-hydride state
on the basis of the spectral and calculated results. The thermodynamics and kinetics of
individual steps of the catalysis were analyzed by density functional theory (DFT) calculations
and electroanalytical simulations. The stepwise CCE or CE (C, chemical; E, electrochemical)
mechanism was operative from the HOTf or TFA source, respectively. In contrast, the
involvement of anilinium acid most likely initiated a proton-coupled electron transfer (PCET) pathway that avoided the
disfavored intermediate after the initial protonation. Via the PCET pathway, the heterogeneous electron transfer rate was
increased and the overpotential was decreased by 0.4 V in comparison with the stepwise pathways. The results showed that the
PCET-involved catalysis exhibited substantial kinetic and thermodynamic advantages in comparison to the stepwise pathway;
thus, an efficient catalytic system for proton reduction was established.
Proton-coupled electron transfer (PCET) is an important
energy conversion process in chemical and biological systems;
it proceeds by coupling protons and electrons together.47−65
The coupled proton−electron transfer occurs by bypassing the
unstable or unfavorable intermediate. Incorporating PCET into
the electrocatalytic reactions could result in energetically and
kinetically more efficient processes.66−73 For electrochemical
PCET, a great deal of effort has been given to the investigation
of the redox behavior of aminophenol compounds, which were
designed to mimic the oxidation of tyrosine in Photosystem II.
In biological systems, histidine nearby acts as a proton acceptor
when a tyrosyl radical is formed. Much progress on the
fundamentals of PCET has been achieved by studying such
aminophenol systems using electrochemical methods74−79 and
theoretical calculations.77−82 From the aspect of hydrogen
evolution, an electrocatalytic system based on a Ni-Ru
molecular catalyst has been reported using Et3NH+ as a proton
source.83 Density functional theory (DFT) calculations
suggested that the PCET reaction occurred. In Co66,84 and
Ni21 hangman porphyrin complexes, intramolecular proton
transfer initiated by the hangman motif facilitated proton
reduction at more positive potentials and led to faster proton
transfer in comparison to that of the analogous non-hangman
systems. Furthermore, a comprehensive theoretical study on a
INTRODUCTION
■
The discovery of new types of inexpensive energy resources is
important for the global economy, even though recent advances
in shale gas extraction technology may provide a temporary
solution to the energy shortage in the near future. To address
this issue and avoid the tremendous environmental impact from
accumulation of greenhouse gases after burnout of fossil fuels,
molecular hydrogen is considered as the next-generation fuel.1,2
Its stored energy (142 MJ/kg) is released accompanied by
water formation, which makes H2 an emission-free energy
vector. Industrial production of hydrogen (>75%) currently
comes from steam re-forming of hydrocarbons.3,4 This process
requires the use of fossil fuels, mainly natural gas, as a feedstock
and elevated working temperatures in the presence of
heterogeneous metal catalysts. This is a well-developed and,
however, heavily energy consuming process that is accom-
panied by the generation of CO2 waste. Electrocatalysis is a
promising technology that converts electric energy into
chemical bonds. Several metal complexes containing earth-
abundant metals such as Fe,5−15 Ni,16−22 Co,23−29 and Mo30,31
have been examined for potential applications in electrocatalytic
hydrogen evolution. For example, electrocatalysts based on the
{Fe2S2} motif have been investigated experimentally32−43 and
theoretically44−46 in the past decade. Catalytic and energetic
efficiencies have seen ongoing advances stemming from the
preparation of novel synthetic complexes or systems.
Received: March 1, 2018
© XXXX American Chemical Society
A
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family of [Ni(PR NR′2)2]2+ complexes for H2 production has
the accessible electrochemical window in CH2Cl2rsedolution. The
corresponding event was clearly observed at E1/2 = −2.37 V
in THF solution (Figure S2a). We reasonably estimated that
the reduction of 1− in CH2Cl2 solution would occur at an
applied cathodic potential of −2.37 V because 1− exhibited
similar oxidation potentials in different media (−0.63 and
−0.62 V in CH2Cl2 and THF solutions, respectively).
In the presence of excess anilinium acid ([PhNH3][BArF24],
pKaMeCN = 10.9),90 the CV of 1− exhibited two irreversible
reduction steps at −1.16 and −1.57 V in CH2Cl2 solution
(Figure 2). The current density of the first reduction increased
2
been performed.85−87 The obtained Pourbaix diagrams and
free-energy profiles provided valuable guidelines for proton
management and ligand design in the PCET process. Recently,
the formation of a Ni/S film during H2-evolving electrocatalysis
was suggested to be initiated by a PCET step.88,89 Despite the
fact that these pioneering works revealed some information on
the kinetic and thermodynamic parameters governing the
catalysis, more studies are necessary for a comprehensive
understanding of the stepwise and concerted pathways in the
electrocatalytic system.
In this work, the detailed mechanism of hydrogen evolution
by [(μ,κ2-bdt)(μ-PPh2)Fe2(CO)5]− in a wide range of acids is
investigated. In contrast to the catalysis in trifluoromethane-
sulfonic acid (HOTf), trichloroacetic acid (TCA), and
trifluoroacetic acid (TFA), which involves stepwise pathways,
the catalytic reaction in the presence of anilinium acid occurs
via a PCET pathway, which facilitates hydrogen evolution at the
same potential as that of the HOTf catalysis. The instant effect
regarding energy usage is that the overpotential is reduced by
+0.4 V. In addition, the heterogeneous electron transfer rate is
enhanced by up to 40 times according to the results of
simulations of electrochemical data. The conjectured stepwise
mechanism in the catalysis of anilinium acid by this catalyst is
ruled out by the results of experiments, DFT calculations, and
electroanalytic simulations. To the best of our knowledge, this
is the first Fe2S2 catalyst which exhibits a PCET-involved
electrocatalysis. Substantial thermodynamic and kinetic benefits
are observed in this effective catalytic system for hydrogen
formation.
Figure 2. Cyclic voltammograms of 1− in CH2Cl2 solution with
increments of anilinium acid (1 mM 1−, υ = 0.1 V s−1, 0.1 M
[nBu4N][BArF24], 295 K) under N2. All potentials were calibrated to
the Fc+/Fc pair as an internal standard. The background CV is shown
as a dashed trace.
RESULTS AND DISCUSSION
■
linearly with the square root of the acid concentration, whereas
the current density of the second reduction did not, indicating
that the first wave was the catalytic process. When the peak
current of the catalytic wave was plotted against the square root
of the scan rate, a linear dependence was obtained. This result
suggests that the catalysis is diffusion-limited and is a
homogeneous reaction.
Redox Chemistry of 1− and Electrocatalytic Hydrogen
Evolution by 1−. Cyclic voltammograms (CVs) of [(μ,κ2-
bdt)(μ-PPh2)Fe2(CO)5]− (1−; bdt = 1,2-benzenedithiolate)
under acid-free conditions showed a reversible one-electron
oxidation event at E1/2 = −0.63 V (vs Fc+/Fc; all potentials
ox
throughout the article are referenced against Fc+/Fc) in
CH2Cl2 solution (Figure 1). This process was diffusion-
controlled, as indicated by the linear dependence of the peak
current on the square root of the scan rate (Figure S2b). A
reduction response was only displayed as the tail at the end of
When electrocatalysis was performed in the presence of
excess TFA (pKaMeCN = 12.7)91 or TCA (pKaMeCN = 10.6),90
one irreversible reduction event at −1.57 V was recorded
(Figure 3b and Figure S9, respectively). The current density
increased linearly with the square root of the acid
concentration, suggesting a catalytic process. Regarding the
electrochemical behavior involving a strong acid, two reduction
waves at −1.16 and −1.57 V were observed for 1− when the
proton source was replaced by HOTf (pKaMeCN = 2.6),90 as
shown in Figure 3a. The peak current of the redox event at
−1.16 V increased linearly with sequential increase of the added
acids (icat ∝ [acid]0.5), suggesting a catalytic response of
hydrogen production. Among these acids, those of medium
strength (TFA and TCA) exhibited one catalytic wave at −1.57
V, whereas the wave was observed at −1.16 V, accompanied by
an additional noncatalytic process at −1.57 V, for a stronger
acid (HOTf). Anilinium acid displayed a different electro-
chemical behavior in comparison to TFA and TCA, even
though they have similar pKa values.
Determination of the Intermediates and Dominant
Reaction Pathways. To elucidate the electrocatalytic hydro-
gen-evolution mechanism for various proton sources, we
characterized intermediate species under acidic conditions
and investigated possible catalytic routes. Reaction of 1− with
Figure 1. Cyclic voltammograms of 1− alone (black), 1μH (from the
reaction of 1− with 1 equiv of HBArF24, blue), and 1μHSH+ (from the
reaction of 1− with 2 equiv of HBArF24, green) in CH2Cl2 solution
(2.4 mM 1−, υ = 0.1 V s−1, 298 K). The background CV is shown as a
dashed trace.
B
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−
a stoichiometric amount of acid (e.g., HBArF24 (BArF24
=
−
B(3,5-C6H3(CF3)2)4 ), HOTf, TCA, anilinium acid, or TFA)
afforded the monoprotonated species [(μ,κ2-bdt)(μ-PPh2)(μ-
H)Fe2(CO)5] (1μH), as depicted in Scheme 1. Additional
amounts of HBArF24 (or HOTf) led to the doubly protonated
species [(μ,κ2-bdtH)(μ-PPh2)(μ-H)Fe2(CO)5]+ (1μHSH+),92
which cannot be obtained in the presence of excess TCA, TFA,
and anilinium acid. The electrochemical behavior of both
protonated species is displayed in Figure 1. Cathodic scans of
HBArF24-yielded 1μHSH+ revealed two irreversible waves at
Epc1(1μHSH+) = −1 V and Epc2(1μHSH+) = −1.57 V and one
irreversible wave at Epc(1μH) = −1.57 V for 1μH. The first
cathodic wave of the HOTf-yielded 1μHSH+ was recorded at
−1.16 V. The shift of 160 mV toward more negative potential
was attributed to the influence of ion pairing, which plays a
non-negligible role in electrochemical responses in solution.93
The weak interaction between the [OTf]− anion and the S
proton of 1μHSH+ partially neutralized the positive charge,
exerting a negative effect relative to 1μHSH+ in the
noncoordinating [BArF24]− salt. Because these cathodic waves
were approximately the same as those for the oxidation pair of
1− with respect to peak height after normalization, every
irreversible reduction event corresponded to a one-electron
charge transfer process when parameters for charge-transfer and
mass-transfer phenomena in the redox reactions were treated in
the same manner.
The chemical reduction of 1μHSH+ by 1 equiv of
= −1.33 V)94 yielded an IR
CH2Cl2
cobaltocene (Cp2Co, E1/2
Figure 3. Cyclic voltammograms of 1− in CH2Cl2 solution with
increments of (a) HOTf and (b) TFA (1 mM 1−, υ = 0.1 V s−1, 0.1 M
[nBu4N][BArF24], 295 K) under N2. All potentials were calibrated to
the Fc+/Fc pair as an internal standard. Background CVs are shown as
dashed traces.
profile that was changed only by a shift of approximately 20
cm−1 to lower energies (Figure S1a). This species was
characterized to be the predecessor 1μH.92 Regeneration of
1−, supported by the IR evidence, was achieved via the chemical
a
Scheme 1. Schematic of Hydrogen Production by 1− in the Presence of Acids
a
Hydrogen evolution occurs to 1μHSH+ and 1μH upon reduction. 1μHSH is a proposed transient species that has not been experimentally
characterized.
C
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reduction of 1μH by decamethylcobaltocene (Cp*2Co,
= −1.94 V)94 (Figure S1b). The results indicate
CH2Cl2
E1/2
that the first reduction in the CV of 1μHSH+ generated 1μH
and that the second wave originated from the reduction of
1μH. The contents of the gaseous products from both redox
events were analyzed by GC and were identified as H2 (Figure
S1c). The reduction processes associated with the protonated
species are summarized in Scheme 1.
The chemical reduction in the presence of acids was also
investigated. 1μH was the essential catalytic species because a
negligible amount of hydrogen (<1% of the H2 amount from
the catalyzed reaction) was detected for the reaction without
the catalyst. Upon the reduction of 1μH to 1μH−, hydrogen
evolution could occur by either or both of the following
mechanisms: a bimolecular reaction that involves loss of H•
radical and a unimolecular reaction that involves the reaction of
hydride with acid. In the former, 1μH− reacted to the other
same species to yield 1− as well as H2, and the as-generated 1−
was readily reprotonated to 1μH. The acid−base reaction via
the unimolecular pathway generated the oxidized product 1 and
H2. These two mechanisms are displayed in Scheme S2.
Characterization of the reaction product enables us to
determine the operating mechanism. When the solution
containing 1 equiv of 1μH and 3 equiv of TFA was subjected
to 1 equiv of Cp*2Co, the solution species characterized by
FTIR spectroscopy was 1μH and acid, accompanied by the
formation of hydrogen bubbles (84% yield, 0.42 equiv of H2).
As the oxidized species 1 and the disproportionation species 1+
were not detected according to the in situ FTIR results, the
possibilities of the heterolytic pathway and the disproportiona-
tion of 1, respectively, were ruled out. The results indicate that
release of H2 from 1μH− is faster than the reaction between
1μH− and acid. The best explanation for the experimental
results is that the bimolecular mechanism is operative for H2
evolution under the limited acid conditions. The same
conclusion was drawn for the 1μHSH+ case.
Figure 4. (a) Cyclic voltammogram of 1μH (1.2 mM) at 50 V s−1 in
CH2Cl2 solution. ipa′ is defined as ipa for the oxidation of 1−. (b) Cyclic
voltammograms of 1μH at 50 V s−1 at [1μH] = 0.6, 1.2, 1.8, 2.4, 3
mM. The inset is the plot of ipa′/ipc vs initial concentration of 1μH (R2
= 0.9901). The background CV is shown as a dashed trace.
(Figure 3a). The proposed H2-evolution pathways were
constructed as follows:
We also used electroanalytical methods to confirm the
bimolecular pathway being the operative catalytic mechanism:
the passage of coulomb charge of the catalysis67,95 and
voltammetric responses as a function of 1μH.96,97 Under the
electroreducing conditions, 1− is produced at the electrode
surface with consumption of one electron per Fe2 if the
bimolecular mechanism occurs. Constant-potential electrolysis
of 1μH at a potential of −1.7 V resulted in the passage of 1.03
equiv of charge and the resultant species was characterized by
FTIR spectroscopy to be 1−. The results suggested a net loss of
H• radical from 1μH− for H2 evolution, following the
reduction. The validity of this route was also evaluated by
monitoring the peak current ratio with varied concentrations of
1μH. The ratio ipa′/ipc reveals a quantitative measure of 1−
formed (ipa′) relative to the concentration of 1μH (ipc). If the
bimolecular route is operative, ipa′/ipc reveals a linear depend-
ence on the initial concentration of 1μH, [1μH]0, in the
absence of acid. As shown in Figure 4, ipa′/ipc increased with
increasing [1μH]0 ranging from 0.6 to 3 mM at υ = 50 V s−1,
indicating that more 1− was generated when the higher [1μH]0
was employed. This lent additional support to the notion that
1μH− decays to 1− and H2 via the dominant homolytic
pathway.
ks′
1μHSH+ + e− HooI 1μHSH
(1)
krxn
′
1μHSH + 1μHSH ⎯⎯⎯⎯→ 1μH + 1μH + H2
Keq′, kPT′
(2)
1μH + HA HoooooooooooooI 1μHSH+ + A
−
(3)
One-electron reduction of the doubly protonated species
1μHSH+ proceeded with electron transfer rate ks′. The
subsequent bimolecular reaction converted 1μHSH into 1μH
with reaction rate krxn′. The catalytic cycle was completed by
the regeneration of 1μHSH+ with proton transfer rate kPT′ and
equilibrium constant Keq′ between 1μH and HOTf. Here, the
bimolecular (i.e., homolytic) pathway (eq 2) was chosen over
the heterolytic pathway (i.e., the hydride-proton reaction) on
the basis of our spectral and electrochemical results (vide
supra). We employed digital simulation of the cyclic
voltammograms on the basis of the proposed mechanism
(Figure S7 and Table S7). The best-fit kPT′ and ks′ values were
2.2 × 105 s−1 M−1 and 0.029 cm s−1, respectively.
Regarding TFA, TCA, and anilinium acid, 1μH was the only
product in the reaction of 1− with excess acids. The irreversible
reduction at −1.57 V was assigned to the 1μH0/− process in the
1−-TFA and 1−-TCA systems. Similar to the catalysis of
1μHSH+, the catalysis of 1μH started with reduction, followed
by hydrogen evolution via the bimolecular pathway and then
Stepwise Mechanism involving HOTf, TCA, and TFA.
The irreversible reduction steps at −1.16 and −1.57 V in the
electrochemical measurements of the 1−-HOTf system were
assigned to the 1μHSH+/0 and 1μH0/− events, respectively
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protonation to complete the cycle. The catalytic mechanism
presumably due to the hydrogen-bond (H-bond) interactions.
The 1H NMR resonance for the H-bonded proton was
observed at 4.54 ppm (Figures S17 and S18). The DFT-
optimized structure of the H-bond-paired species
1μHSHNH2Ph+ revealed a short S···H contact of 2.135 Å
from one acidic proton of the anilinium cation toward the
terminal thiolate of 1μH (Figure S16c). This distance remained
well within the sum of van der Waals radii. A similar H-bonding
interaction between the St of 1− and N,N-dimethylanilinium
acid was recently characterized at 183 K.100 This observed H-
bonding interaction supports the hypothesis that the St site of
1μH acts as the proton acceptor during catalysis. We then
considered that the reduction accounting for the other
mechanism was associated with the H-bonding interaction.
The resulting H-bonded species 1μHSHNH2Ph+ was more
stable than 1μH by approximately 5 kcal mol−1 but was still
unfavorable. A much smaller energy shift was anticipated for the
transition from the 1μH state to the 1μHSHNH2Ph+ state than
for the conversion from 1μH to 1μHSH+ because the H-
bonding interaction has less influence on the electronic
structure of the complex in comparison to the full protonation.
The calculated potential Ecal(1μHSHNH2Ph+/0) was −1.69 V.
The resulting value showed a large deviation from the −1.16 V
obtained for the PCET process, even after a calibration of 220
mV was used to account for the theory−experiment difference
for E(1μHSH+/0). This result indicates that the driving force
from the H-bonding interaction was not sufficiently strong to
give rise to the positive potential shift that was compatible with
the direct protonation. These results lent support to the
suggestion that the PCET catalysis occurred in the 1μH-
concerning 1μH0/− was described as follows:
ks″
1μH+ + e‐ HooI 1μH−
(4)
krxn
″
1μH− + 1μH− ⎯⎯⎯⎯⎯→ 1− + 1− + H2
Keq″, kPT″
(5)
1− + HA HooooooooooooooI 1μH + A
−
(6)
The corresponding rate constant for each step was obtained
from the least-squares fits to the experimental results, giving
kPT″ and ks″ values of 3495 s−1 M−1 and 0.0162 cm s−1,
respectively, for the TFA system. For the TCA system, the
corresponding parameters were 4198 s−1 M−1 and 0.0219 cm
s−1, respectively (Figures S8 and S10 and Tables S8 and S9).
In contrast to TFA and TCA, the different electrochemical
behavior observed for anilinium acid suggested a new catalytic
mechanism. Because the potentials for the catalytic responses of
the 1μH-[PhNH3][BArF24] system and 1μHSH+ were very
similar, we proposed an intermediate similar to the 1μHSH
+
species involved in the 1μH-PhNH3 system.
Elucidation of the PCET Event. Because both of the
individual steps, i.e., proton transfer (PT) and electron transfer
(ET), were not initially spontaneous in the reaction scheme of
1μH in the 1μH-PhNH3+ system, the catalytic process at −1.16
V would most likely occur via a PCET step. We performed
DFT calculations to obtain insights into the mechanistic routes
and intermediates. Figure 5 gives a schematic of the states and
+
PhNH3 system.
Kinetic Study of the PCET Event. To disentangle the
kinetic behavior of this PCET event (1μHSHPCET), we used a
nonadiabatic Marcusian electron transfer formula for PCET
step in the digital simulation. The mechanism was depicted as
follows:
kPCET
1μH + PhNH3+ + e− HooooooI 1μHSH + PhNH2
(7)
krxn
′
1μHSH + 1μHSH ⎯⎯⎯⎯→ 1μH + 1μH + H2
(8)
Figure 5. Theoretical thermochemistry of the 1μH-PhNH3+ system in
the catalysis. The dashed lines indicate the disfavored pathways. The
calculated redox potentials are referenced to the Fc+/Fc pair.
The cyclic voltammograms of 1μH in the presence of
stoichiometric amounts of PhNH3 were obtained. The
+
apparent heterogeneous rate constant kPCET was optimized by
fitting the cyclic voltammograms from different scan rates and
various acid concentrations (Figure S11 and Table S10). The
species 1μHSH converts to 1μH at a hydrogen production rate
krxn′ assumed to be 1.92 × 1010 s−1 M−1, which is the value
taken from the HOTf catalysis. The best-fitted results gave 0.63
cm s−1 for the kPCET value.
It has also been pointed out that the potential shift between
the catalytic and noncatalytic events may occur in stepwise
reactions.76,84,101,102 To examine the validity of the stepwise
mechanism, we therefore decoupled the PCET event into
hypothetical ETPT pathways as follows:
+
charge transfer reactions in the 1μH-PhNH3 catalytic system.
In the first stepwise mechanism, the initial PT was
thermodynamically disfavored, with ΔG°PT = 10.78 kcal
mol−1. This result is consistent with the absence of 1μHSH+
in solution. The subsequent ET occurred at the potential
Ecal(1μHSH+/0) = −1.22 V. This step was observed in the
HOTf catalysis by 1−. This calculated redox potential value was
considered to be reliable because any potential difference
between the calculated and the experimental results was
acceptable if it was smaller than 400 mV.98,99
Furthermore, we observed that, when 1μH was in contact
with excess anilinium acids (>20 equiv) at ambient temper-
ature, the IR profile of 1μH in solution shifted to higher
energies by +8 cm−1, with the exhibition of three distinct νCO
bands at 2088 (s), 2035 (vs), and 1985 (m) cm−1 and one
shoulder at 2022 cm−1 (Figure S16a). The moderate energy
difference suggested that the electronic structure of 1μH was
partially perturbed by the surrounding PhNH3+ proton donors,
ks″
1μH + e− HooI 1μH−
(9)
kPT ′′′
1μH− + PhNH3+ HoooooooI 1μHSH + PhNH2
(10)
krxn
′
1μHSH + 1μHSH ⎯⎯⎯⎯→ 1μH + 1μH + H2
(11)
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Table 1. Summary of Electrocatalytic Results
a
b
c
c
acid
CF3SO3H
pKaMeCN
Ecat (V)
overpotential (V)
kPT (s−1 M−1
2.2 × 105
4198
)
ks (cm s−1
)
process
2.6
10.6
12.7
10.7
−1.16
−1.57
−1.57
−1.16
+0.74
+0.86
+0.82
+0.40
0.0290
0.0219
0.0162
0.63
1μHSH+/0
1μH0/−
1μH0/−
1μHSHPCET
CCl3COOH
CF3COOH
3495
[PhNH3][BArF24]
N/A
a
b
All potentials are reported against the Fc+/Fc couple. Overpotential = |EH+ − Ecat/2|. The thermodynamic potential (EH+) has been calibrated to a
value for 10 mM of acid on the basis of eq 8 in Inorg. Chem. 2010, 49, 10338, except for HOTf, which is referenced to Inorg. Chem. 2006, 45, 9181.
c
kPT and ks in the table are the general descriptions for individual PT and ET rates, respectively, in the different pathways.
Upon one-electron reduction, proton transfer simultaneously
occurred at the rate kPT′′′ to generate 1μHSH as the ready-
state species for H2 evolution. The intramolecular pathway of
H2 evolution (1μHSH → 1 + H2) was ruled out according to
the DFT results (Scheme S1). In the simulations, the diffusion
coefficients and krxn′ values were assumed to be the same as
those from the PCET pathways. The proton transfer rate kPT′′′
was initially assumed to be the maximum diffusion rate kMax‑diff
substantial thermodynamic benefit for the catalysis. Impor-
tantly, the H/D kinetic isotope experiments showed similar
catalytic current responses and identical reduction potentials
under the same conditions (Table S11). The results suggest
that the barrier in the intersection of the proton potential
energy curves of the electrochemical PCET is small and that
the overlap in the ground proton vibronic states is large. The
PCET also affected the heterogeneous ET rate. The rate for the
1μHSHPCET process was on the order of 0.63 cm s−1, which is
up to 40 times greater than the estimated rate of 0.016−0.029
cm s−1 for the ordinary 1μH0/− and 1μHSH+/0 processes.
Overall, a decrease in the working potential and enhancement
of the heterogeneous ET rate in the PCET event lead to an
improvement in the catalytic performance relative to that of the
stepwise processes (Table 1).
+
between 1μH− and PhNH3 , which was estimated by the
Debye−Smoluchowski method89,103 (see the Supporting
Information). The transfer coefficient α and ET rate ks′′ were
separately tuned in curve fitting until the best fit on either the
current amplitude or catalytic peak potential was achieved.
Neither the experimental peak potential shift nor the current
amplitude was reproduced in the simulations with the
application of the kMax‑diff value of 7.8 × 109 s−1 M−1 (Figure
S13a−d). With an increase in the ks′′ rate, the simulated peak
potential anodically shifted to a potential boundary, ∼− 1.35 V.
The results suggest the intermolecular PT to be a rate-limiting
step in this hypothetical stepwise mechanism. In the previous
estimation of kMax‑diff, the electrostatic interaction between the
CONCLUSION
■
In summary, the redox behavior of the Fe−S-containing catalyst
1− was investigated and its electrocatalytic properties under
various acidic conditions were characterized to proceed via two
different hydrogen evolution processes that were ascribed to
stepwise CCE (C, chemical reaction; E, one-electron
reduction) and CE processes concerning the Fe-hydride-S-
proton species and the Fe-hydride species, respectively. The use
of anilinium acid as the proton source allowed the catalysis to
occur via the PCET route. This pathway is associated with the
concomitant PT from PhNH3+ to the St of 1μH, along with ET
from the electrode surface to 1μH. Here, we presented an
example of the PCET catalysis for H2 formation and showed
that the careful management of Brønsted acid−base couples is
capable of enhancing the catalytic performance. Further study
of the electrocatalytic mechanisms associated with the proton-
donating/-accepting ability is ongoing in our laboratory.
two charged species, 1μH− and PhNH3 , was not considered.
+
Corr
Max−diff
The maximum diffusion rate was corrected to k
as the
electrostatic interaction was included. Upon substituting a
Corr
Max−diff
k
value of 2.15 × 1015 s−1 M−1 for kMax‑diff, the best fit to
the peak potential was able to be achieved by varying α or ks′′.
The resultant current amplitude, however, was largely deviated
from the experimental data (Figure S13e−h). We could not
obtain a satisfactory fit that fully reproduced the experimental
current amplitude and catalytic peak potential when optimizing
α and ks′′ in the simulation. It is said that the catalysis in the
+
1μH-PhNH3 system cannot be described as a stepwise ETPT
reaction.
Discussion of the PCET vs the Non-PCET Event. In
these studies, the catalytic performance of 1− under various
acids with pKa values between 2.6 and 12.7 was investigated.
Two observed catalytic processes were due to the 1μHSH+/0
and 1μH0/− processes under the conditions of strong- and
medium-strength acids, respectively. Notably, the employment
of anilinium acid facilitated the PCET-involved catalytic
process. To elucidate whether the pKa value of an acid was
decisive for the PCET reaction in our system, we compared the
catalysis of TCA and anilinium acid, where both acids have
similar pKa values and afforded only 1μH in solution. The
EXPERIMENTAL SECTION
■
General Methods. All reactions were carried out by using standard
Schlenk and vacuum-line techniques under an atmosphere of purified
nitrogen. All commercially available chemicals from Aldrich were of
ACS grade and were used without further purification. Solvents were
of HPLC grade and were purified as follows. Hexane, diethyl ether,
and tetrahydrofuran (THF) were distilled from sodium/benzophe-
none under N2. Dichloromethane was distilled from CaH2 under N2.
Acetonitrile (MeCN) was distilled first over CaH2 and then from P2O5
under N2. Deuterated solvents obtained from Merck were distilled
over 4 Å molecular sieves under N2 prior to use.
+
electrocatalytic event occurred at −1.16 V for the 1−-PhNH3
Infrared spectra were recorded on a PerkinElmer Spectrum One
instrument using a 0.05 mm CaF2 cell. H, 13C{1H}, and 31P{1H}
1
system, more positive by 0.4 V than for the 1−-TCA system.
These results suggest that cationic anilinium acid rather than
the pKa of the proton sources was indispensable for the PCET
reaction. For the overpotential, an energy savings of 0.4 V
relative to the conventional stepwise processes (+0.4 V vs
approximately +0.8 V, Table 1) was observed when the PCET
process occurred, indicating that the PCET reaction offers a
NMR spectra were recorded on a Bruker AV-500 or DRX-500
spectrometer operating at 500, 125.7, and 202.49 MHz, respectively.
1H-DOSY NMR spectra were recorded on a 600 MHz Varian VNMRS
spectrometer at 298 K. Mass spectral analyses were done on a Waters
LCT Premier XE instrument at the Mass Spectrometry Center of the
Institute of Chemistry, Academia Sinica. Elemental analyses were
performed on an Elementar vario EL III elemental analyzer. Analysis
F
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Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
of H2 gas from controlled-potential electrolysis and chemical reduction
reactions was performed on an Agilent 7890 gas chromatograph
equipped with a thermal conductivity detector (TCD) and fitted with
a Restek ShinCarbon ST column (100/120 mesh, 2 m, 1/16 in. o.d.,
1.0 mm i.d.). The column temperature was kept at 80 °C, and the
injection and detector were both at 200 °C. The gas flow was 15 mL
min−1. Nitrogen was used as the carrier gas.
Molecular Structure Determination. The X-ray single-crystal
crystallographic data collection for [TBA][1] (CCDC 976229) was
carried out at 150 K on a Bruker SMART APEX CCD four-circle
diffractometer with graphite-monochromated Mo Kα radiation (λ =
0.71073 Å) outfitted with a low-temperature, nitrogen-stream aperture.
The structures were solved using direct methods, in conjunction with
standard difference Fourier techniques, and refined by full-matrix least-
squares procedures. A summary of the crystallographic data for
complex 1− is shown in Table S1. An empirical absorption correction
(multiscan) was applied to the diffraction data for all structures. All
non-hydrogen atoms were refined anisotropically, and all hydrogen
atoms were placed in geometrically calculated positions by the riding
model. All software used for diffraction data processing and crystal
structure solution and refinement are contained in the SHELXTL97
program suites.104 In [TBA][(μ,κ2-bdt)(μ-PPh2)Fe2(CO)5]·THF, two
of the butyl carbon atoms in the nBu4N+ cation (C34, C35) are
disordered.
the daughter species [(μ-bdt)Fe2(CO)5(PPh2H)] and [(μ,μ,κ2-
bdt)4(μ-PPh2)4Fe8(CO)18]. 1μH was reported to slowly convert to
[(μ-bdt)Fe2(CO)5(PPh2H)] in solution.92 It was reported that
[(μ,μ,κ2-bdt)4(μ-PPh2)4Fe8(CO)18] was generated from oxidation of
1−.106
Synthesis. Synthesis of [TBA][(μ,κ2-bdt)(μ-PPh2)Fe2(CO)5] ([TBA]-
[1]). The preparation procedures are similar to those published
previously.92 To a red THF solution (30 mL) of [(μ-bdt)-
Fe2(CO)6]107 (500 mg, 1.19 mmol) was added a 0.5 M solution of
KPPh2 (2.4 mL, 1.2 mmol) in THF. The brown solution was stirred
overnight and dried in vacuo. The green-brown solution was mixed
with 50 mL of acetone, followed by addition of 330 mg (1.19 mmol)
of tetrabutylammonium chloride (TBACl). The reaction mixture was
stirred at room temperature for 3 h, and the solvent was removed in
vacuo. The residue was dissolved in 20 mL of CH2Cl2, and the
solution was filtered through Celite to remove insoluble salt. The
product was precipitated upon addition of hexane. The green-brown
solid was washed three times with 20 mL of ether and 30 mL of
hexane and then dried under vacuum. The yield was 761 mg (78%).
Crystals of [TBA][(μ,κ2-bdt)(μ-PPh2)Fe2(CO)5]·THF suitable for X-
ray crystallographic analysis were grown from a THF/hexane solution
at −20 °C. IR (THF, cm−1): νCO 2005 m, 1961 vs, 1936 m, 1916 s,
1
1896 m. H NMR (500 MHz, d-THF, 295 K): 0.97 (t, 12H, CH3 of
TBA+), 1.36 (m, 8H, CH2 of TBA+), 1.60 (m, 8H, CH2 of TBA+), 3.15
(m, 8H, NCH2 of TBA+), 6.63 (t, 1H, 3JHH = 7.0 Hz, S2C6H4), 6.72 (t,
Electrochemistry. Electrochemical measurements were recorded
on a CH Instruments 630C electrochemical potentiostat using a
gastight three-electrode cell under N2 at 298 K unless otherwise stated.
A vitreous carbon electrode (1 or 3 mm in diameter) and a platinum
wire were used as working and auxiliary electrodes, respectively. The
3
1H, JHH = 7.0 Hz, S2C63H4), 7.11 (m, 3H, P(C6H5)2), 7.21 (m, 3H,
P(C6H5)2), 7.42 (d, 1H, JHH = 8 Hz, S2C6H4), 7.54 (m, 3H, S2C6H4,
P(C6H5)2), 7.72 (m, 2H, P(C6H5)2) ppm. 13C{1H} NMR (125.7
MHz, d-THF, 298 K): 12.95 (s, 4C, CH3 of TBA+), 19.66 (s, 4C, CH2
of TBA+), 23.68 (s, 4C, CH2 of TBA+), 58.49 (s, 4C, NCH2 of TBA+),
reference electrode was a nonaqueous Ag+/Ag electrode (0.01 M
−
AgNO3/0.1
M
[nBu4N][BArF24],105 BArF24
= B(3,5-
117.86 (s, 1C, S2C6H4), 122.97 (s, 1C, S2C6H4), 126.61 (d, 2C, JPC
=
−
C6H3(CF3)2)4 ). The reference electrode was in a separate compart-
ment with a Vycor glass tip. The working and counter electrodes were
in the same compartment when the CVs were collected. The working
electrode was polished after each scan. All potentials were measured in
CH2Cl2 or THF solution containing 0.1 M [nBu4N][BArF24]. The
potentials are reported against the ferrocenium/ferrocene (Fc+/Fc)
couple. All measurements were performed with iR compensation. The
simulation of the cyclic voltammetry was performed using the DigiElch
7.FD program on the background-corrected data.
9.3 Hz, P(C6H5)2), 127.16 (d, 2C, JPC = 9.7 Hz, P(C6H5)2), 127.41 (s,
1C, P(C6H5)2), 127.91 (s, 1C, P(C6H5)2), 128.76 (s, 1C, S2C6H4),
129.33 (s, 1C, S2C6H4), 133.68 (m, 4C, P(C6H5)2), 139.59 (d, 1C, J =
21.3 Hz, ipso-S2C6H4), 141.25 (d, 1C, JPC = 34.7 Hz, ipso-P(C6H5)2),
144.09 (d, 1C, JPC = 26.3 Hz, ipso-P(C6H5)2), 160.17 (s, 1C, ipso-
S2C6H4), 216.67 (d, 3C, JPC = 3.4 Hz, CO), 218.81 (d, 1C, JPC = 10.2
Hz, CO), 220.98 (d, 1C, JPC = 18.9 Hz, CO) ppm. 31P{1H} NMR
(202.48 MHz, d-THF, 296 K): 144.7 (s) ppm. ESI-MS: m/z 576.9
{M}−, 548.9 {M − CO}−, 520.9 {M − 2CO}−, 492.9 {M − 3CO}−,
464.9 {M − 4CO}−, 436.9 {M − 5CO}−. Anal. Calcd for
C39H50Fe2NO5PS2: N, 1.71; C, 57.15; H, 6.15. Found: N, 1.92, C,
57.04; H, 6.12.
Controlled-potential electrolysis (CPE) was performed in a three-
electrode cell under N2 at 298 or 273 K. The working and counter
electrodes were graphite rods (6.15 mm in diameter) and were
separated in two compartments when the CPE experiments were
performed. The compartment containing the counter electrode had a
10 mm diameter fine-porosity glass frit. The CPEs were performed at
−1.1 and −1.6 V (vs Fc+/Fc) for hydrogen production from 1μHSH+
and 1μH, respectively. The potential for the PCET reaction involving
1μH in the presence of anilinium acid was performed at −1.1 V (vs
Fc+/Fc). Regarding the Faradaic efficiency (FE) and turnover number
(TON), three runs were conducted at each potential for 1 h at 298 K.
The FTIR results indicated that the catalyst remained stable during the
CPE conducted at 273 K. The reported FEs and TONs were the
average values over three sets of the data. Following CPE, 250 μL of
headspace gas was removed from the working electrode compartment
with a gastight syringe (Hamilton SampleLock) and was analyzed on
an Agilent 7890 gas chromatograph. The amount of hydrogen was
determined in comparison to a standard calibration curve from known
amounts of hydrogen. From the determined amount of hydrogen
generated, TON and FE were calculated according to the following
equations, respectively: TON = n/ncat; FE (%) = 2nF/Qcat. × 100%,
where n is the number of moles of hydrogen produced, ncat is the
number of moles of the catalyst, F is the Faraday constant, and Qcat is
the total amount of charge passed in the CPE.
Reaction of [TBA][(μ,κ2-bdt)(μ-PPh2)Fe2(CO)5] with [PhNH3]-
[BArF24]. To a CH2Cl2 solution (10 mL) of [TBA][(μ,κ2-bdt)(μ-
PPh2)Fe2(CO)5] (100 mg, 0.122 mmol) was added a CH2Cl2 solution
(10 mL) of [PhNH3][BArF24]108 (143 mg, 0.15 mmol). The IR profile
was changed from 2005 m, 1961 vs, 1936, 1916, and 1896 m cm−1 to
2080 s, 2030 vs, 2013 s and 1983 m cm−1, as monitored by FTIR
spectroscopy. The profile was shifted to higher energy upon addition
of an additional 29 equiv of [PhNH3][BArF24] in CH2Cl2 solution. IR
(CH2Cl2, cm−1): νCO 2088 s, 2035 vs, 2022 sh, 1985 m cm−1.
Reaction of [TBA][(μ,κ2-bdt)(μ-PPh2)Fe2(CO)5] with CF3COOH. A
CH2Cl2 solution (5 mL) of [TBA][(μ,κ2-bdt)(μ-PPh2)Fe2(CO)5] (50
mg, 0.061 mmol) was treated with CF3COOH (10 μL, 0.13 mmol).
The IR profile was changed from 2005 m, 1961 vs, 1936, 1916, and
1896 m cm−1 to 2081 s, 2030 vs, 2013 s, and 1983 m cm−1, as
monitored by FTIR spectroscopy. The profile was shifted to higher
energy upon addition of an further 28 equiv of CF3COOH in CH2Cl2
solution. IR (CH2Cl2, cm−1): νCO 2083 s, 2034 vs, 2022 sh, 1985 m
cm−1.
Reaction of [TBA][(μ,κ2-bdt)(μ-PPh2)Fe2(CO)5] with CCl3COOH. A
CH2Cl2 solution (5 mL) of [TBA][(μ,κ2-bdt)(μ-PPh2)Fe2(CO)5] (50
mg, 0.061 mmol) was treated with CCl3COOH (13 mg, 0.08 mmol).
The IR profile was changed from 2005 m, 1961 vs, 1936, 1916, and
1896 m cm−1 to 2081 s, 2031 vs, 2013 s, and 1984 m cm−1, as
monitored by FTIR spectroscopy. The profile was unchanged upon
addition of a further 29 equiv of CCl3COOH in CH2Cl2 solution.
Prolonged CPE concerning anilinium acid was performed at 298 K
for 12 h, in which the catalytic current revealed a significant decrease
after t > 7000 s (see the Supporting Information). The major solution
species was 1μH according to the FTIR results. Most of the proton
source was consumed, which is possibly responsible for the current
decay over time. The other plausible explanation is the formation of
Reduction of [(μ,κ2-bdtH)(μ-PPh2)(μ-H)Fe2(CO)5][OTf] ([1μHSH]-
92
[OTf]) by Cobaltocene (Cp2Co). Chemical reduction of 1μHSH+
G
DOI: 10.1021/acs.inorgchem.8b00543
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
was conducted in CH2Cl2 solution at 299 K. To a CH2Cl2 solution (7
mL) of 1μHSH+ (100 mg, 0.137 mmol) was added a CH2Cl2 solution
(5 mL) of cobaltocene (26 mg, 0.137 mmol). An instant color change
to green was observed. The original IR bands were shifted from 2101
s, 2059 vs, 2045 s, and 2019 m cm−1 to lower energies at 2081 s, 2030
vs, 2013 s, and 1984 m cm−1, confirming the formation of 1μH. The
content of gaseous products in the headspace of the Schlenk flask was
analyzed by GC to be H2.
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ASSOCIATED CONTENT
■
S
* Supporting Information
The Supporting Information is available free of charge on the
ACS Publications website at DOI: 10.1021/acs.inorg-
chem.8b00543.
Results of spectroscopy, electrochemistry, gas chroma-
tography, and computational chemistry (PDF)
Accession Codes
CCDC 976229 contains the supplementary crystallographic
data for this paper. These data can be obtained free of charge
via www.ccdc.cam.ac.uk/data_request/cif, or by emailing data_
request@ccdc.cam.ac.uk, or by contacting The Cambridge
Crystallographic Data Centre, 12 Union Road, Cambridge CB2
1EZ, UK; fax: +44 1223 336033.
AUTHOR INFORMATION
■
Corresponding Authors
*E-mail for Y.-C.L.: yuchiao@gate.sinica.edu.tw.
*E-mail for M.-H.C.: mhchiang@chem.sinica.edu.tw.
ORCID
Ming-Hsi Chiang: 0000-0002-7632-9369
Notes
The authors declare no competing financial interest.
(15) Nguyen, A. D.; Rail, M. D.; Shanmugam, M.; Fettinger, J. C.;
Berben, L. A. Electrocatalytic Hydrogen Evolution from Water by a
Series of Iron Carbonyl Clusters. Inorg. Chem. 2013, 52, 12847−
12854.
ACKNOWLEDGMENTS
■
(16) Han, Z. J.; McNamara, W. R.; Eum, M. S.; Holland, P. L.;
Eisenberg, R. A nickel thiolate catalyst for the long-lived photocatalytic
production of hydrogen in a noble-metal-free system. Angew. Chem.,
Int. Ed. 2012, 51, 1667−1670.
This work was supported by the Ministry of Science and
Technology of Taiwan and Academia Sinica. Professor
Marcetta Y. Darensbourg (TAMU) and Professor Wen-Feng
Liaw (NTHU) are acknowledged for helpful discussions. The
authors are grateful to Professor Thomas B. Rauchfuss (UIUC)
for his insightful suggestions. The authors also thank Dr. Mei-
Chun Tseng (AS), Ms. Mei-Ying Chung (AS), and Ms. Li-
(17) Le Goff, A.; Artero, V.; Jousselme, B.; Tran, P. D.; Guillet, N.;
́ ́
Metaye, R.; Fihri, A.; Palacin, S.; Fontecave, M. From hydrogenases to
noble metal-free catalytic nanomaterials for H2 production and uptake.
Science 2009, 326, 1384−1387.
(18) Stewart, M. P.; Ho, M.-H.; Wiese, S.; Lindstrom, M. L.;
Thogerson, C. E.; Raugei, S.; Bullock, R. M.; Helm, M. L. High
Catalytic Rates for Hydrogen Production Using Nickel Electrocatalysts
with Seven-Membered Cyclic Diphosphine Ligands Containing One
Pendant Amine. J. Am. Chem. Soc. 2013, 135, 6033−6046.
(19) Reback, M. L.; Buchko, G. W.; Kier, B. L.; Ginovska-Pangovska,
B.; Xiong, Y.; Lense, S.; Hou, J.; Roberts, J. A. S.; Sorensen, C. M.;
Raugei, S.; Squier, T. C.; Shaw, W. J. Enzyme Design from the Bottom
Up: An Active Nickel Electrocatalyst with a Structured Peptide Outer
Coordination Sphere. Chem. - Eur. J. 2014, 20, 1510−1514.
(20) Helm, M. L.; Stewart, M. P.; Bullock, R. M.; DuBois, M. R.;
DuBois, D. L. A synthetic nickel electrocatalyst with a turnover
1
Ching Shen (NCTU) for assistance with MS, EA, and H-
DOSY NMR measurements, respectively. We are grateful to the
High Performance Computing Service of Academia Sinica
Computer Center (ASCC) and the National Center for High-
performance Computing (NCHC) for computer time and
facilities. The authors also thank the reviewers for helpful
suggestions on mechanistic elucidation.
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