Electrochemical and Theoretical Studies on Nickel

九州大学学術情報リポジトリKyushu University Institutional Repository

Electrochemical and Theoretical Studies onNickel Dithiolene Hydrogen Evolution Catalysts:Developing Ligand-based Proton-coupled ElectronTransfer Pathways

小柴, 慧太

http://hdl.handle.net/2324/2236028

出版情報：九州大学, 2018, 博士（理学）, 課程博士バージョン：権利関係：

Electrochemical and Theoretical Studies on

Nickel Dithiolene

Hydrogen Evolution Catalysts:

Developing Ligand-based

Proton-coupled Electron Transfer Pathways

Keita Koshiba

March 2019

Department of Chemistry

Graduate School of Science

Kyushu University

Contents

General Introduction 1

Hydrogen Evolution Reaction (HER) 1

HER in Nature: [NiFe] and [FeFe] hydrogenases 2

Molecular Catalysis for Electrochemical HER 5

Ligand-based PCET Reduction for HER 10

HER Catalyzed by Bis(dithiolene) Complexes 15

Mechanism of HER Catalyzed by Bis(dithiolene) Complexes 16

Survey of This Thesis 19

References 20

Chapter 1: A Nickel Dithiolate Water Reduction Catalyst Providing Ligand-based

Proton-coupled Electron Transfer Pathways 23

Introduction 23

Experimental Section 26

Results and Discussion 28

Conclusions 43

References 45

Chapter 2: Consecutive Ligand-based PCET Processes Affording a Doubly Reduced

Nickel Pyrazinedithiolate which Transforms into a Metal Hydride Required to

Evolve H2 48

Introduction 48



Conclusions 72

References 74

Chapter 3: Ligand-based PCET Reduction in a Heteroleptic Ni(bpy)(dithiolene)

Electrocatalyst Leading to a Lower Overpotential for Hydrogen Evolution 77

Introduction 77



Conclusions 115

References 116

Concluding Remarks 119

Acknowledgements 121

List of Publications 123

Other Publications 124

1

General Introduction

Hydrogen Evolution Reaction (HER) In last decades, the development of a direct energy conversion system using

renewable energies (e.g., solar energy, hydropower, or wind power) to make chemical

energies (e.g., hydrogen, alchol, or other carbonhydrates) has attracted much attention

due to the limitation of fossil fuels and the impending global warming. Until now,

effective conversion of renewable energies into an electricity has been thus far achieved,

however, effective conversion of an electricity into chemical energies still remains

immature.1-3 In this context, hydrogen evolution reaction (HER) using an electricity has

been extensively studied because HER is considered as one of the most ideal

methodologies to achieve an effective conversion of an electricity into chemical energies.

In order to achieve efficient electrochemical HER from water, various efforts have been

thus far made to develop highly efficient and robust catalysts for HER,1,2 and some rare

metals, such as platinum, rhodium, palladium and so on, have been proved to show

efficient catalytic performance4 although the extremely small abundance of these rare

metals is an intrinsic limitation for the practical application of these catalysts.

Consequently, there is a strong demand to develop highly efficient and robust catalysts

using earth-abundant elements (materials).

In this context, recently, extensive studies on the development of homogeneous

electrocatalysts (i.e., molecular catalysts) for HER have been made. One of the

significances of studying homogeneous catalysts is clarifing, mimicking and re-creating

natural enzymes for HER, which consist of coordination compounds of earth-abundant

elements such as iron(Fe) and nickel(Ni). Revealing natural systems provides important

strategies towards the development of highly effective earth-abundant catalysts for

hydrogen evolution.

2

HER in Nature: [NiFe] and [FeFe] hydrogenases Nature developed efficient enzymatic systems for HER. [NiFe] and [FeFe]

hydrogenases (Scheme 1)5 can promote hydrogen evolution and oxidation reactions

reversibly and effectively with low overpotential (~ 0 mV), and high turnover frequency

(700 s-1 and 6000-21000 s-1, respectively),5a,6 although they have only 1st-row metal ions.

They are expected as one of the alternative catalysts for HER, thus the mechanisms of

their catalysis have been widely studied.

Scheme 1. Reaction centers of [NiFe] and [FeFe] hydrogenases.

The proposed mechanism of HER by [NiFe] hydrogenase is shown in Figure 1.5c,d

For HER, the Ni(II) center is initially reduced to a Ni(I) species coupled with protonation

on the thiolate of terminal cysteine based on a proton-coupled electron transfer (PCET)

process,7 forming a Ni-L state. The second step is an intramolecular proton transfer from

sulfur to dinuclear metal centers to form the hydrido-bridged Ni(III)Fe(II) intermediate,

which is described as a Ni-C state. It can be further reduced to a Ni(II)-H-Fe(II) species,

whose process is coupled with the proton transfer to the thiolate forming a Ni-R state,

followed by releasing hydrogen. In this catalytic cycle, the mechanism can be

summarized as an ECEC pathway, where E is the electrochemical step, and C is the

chemical step (i.e. protonation in the case for HER),8 via consecutive metal-based PCET

reductions. It can be concluded that the cysteine residue and the nickel-iron bridge can

work as the proton acceptor.

[NiFe] hydrogenase [FeFe] hydrogenase

3

Figure 1. Reaction mechanism of HER catalyzed by [NiFe] hydrogenase.

4

The mechanism of HER by [FeFe] hydrogenase is depicted in Figure 2.5c,9 In this

system, an aza-dithiolate ligand, which bridges two iron atoms and the amine axially

located above one of the iron centers, can work as the proton relay. It is realized that the

[FeFe] hydrogenase stabilizes the terminal hydride, which is in sharp contrast with the

bridging hydride intermediates observed in Ni-C and Ni-R states of [NiFe] hydrogenase.

Furthermore, the [4Fe-4S] subcluster tethered to one of the two iron ions serves as an

electron relay, which supports the iron centers undergoing the hydride formation and

hydrogen elimination.

Figure 2. Reaction mechanism of HER catalyzed by [FeFe] hydrogenase.

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5

Molecular Catalysis for Electrochemical HER The HER promoted by molecular catalysts has been studied in the last half century.

Its mechanism was analyzed by several approaches, especially by the electrochemical

analysis using cyclic voltammetry (CV), whose theory was mostly established by Savéant

and co-workers. In this section, the typical mechanisms of electrochemical HER by

molecular catalysts are briefly discussed with referring to Savéant’s works.8

HER comprises of two electrons and two protons transfer processes, where an

electron transfer step and a protonation step, including a hydrogen-evolving step, are

described as ET or E, and PT or C, respectively. For example, the mechanism of HER by

[FeFe] hydrogenase (as discussed above) can be explained as an ET-PT-ET-PT path or

ECEC path. It should be noted here that the HER by hydrogenases and some of molecular

catalysts proceeds via the heterolytic path, in which hydrogen evolves with the reaction

of hydride (H-) and proton (H+). Scheme 2 shows two typical mechanisms of HER by

molecular catalysts (i.e., heterolytic EECC (Scheme 2a) and ECEC (Scheme 2b)

mechanisms). For the heterolytic EECC mechanism, the second reduction triggers the

hydride formation followed by hydrogen evolution. The second reduction is often coupled

with a protonation, which is ascribable as PCET (Scheme 2a). On the other hand,

heterolytic ECEC undergoes the hydride formation process coupled with the first

reduction process (i.e., also described as PCET) (Scheme 2b).

Scheme 2. Square schemes for HER based on EECC mechanism (a) and ECEC mechanism (b).

Mn M(n-1) M(n-2)

M(n+1)(-H) Mn(-H)

2Mn + H2 Mn + H2

C

E

+e-

+H+ C+H+

CC+M(n+1)(-H)

homolysis

E

+e-Mn M(n-1) M(n-2)

M(n+1)(-H) Mn(-H)

2Mn + H2 Mn + H2

C

E

+e-

+H+ C+H+

CC+M(n+1)(-H)

homolysis

E

+e-

(b) Heterolytic ECEC mechanism(a) Heterolytic EECC mechanism

+H+

heterolysis+H+

heterolysis

E

+e-

E

+e-

6

In order to evaluate the catalytic activities of molecular catalysts, a turnover

frequency (TOFmax [s-1]), which relates to a catalytic rate (�� [s-1]) and some rate

constants (i.e., �� , �� and � [s-1 M-1�� and overpotential

(η [V]) are employed. It is desirable if the catalysts show a high catalytic rate with a low

overpotential. The catalytic rate (i.e., ��) and the global rate constant of the catalysis

(i.e., ��) can be obtained by the acid concentration dependence of catalytic currents

(Figure 3) as in eqs. (1) and (2),

�� = 20.446�� 1� �� = �� 2�

where icat denotes to the catalytic peak current, ip is the one-electron reduction peak

current in the absense of acid, R [C V mol-1 K-1] is the gas constant, T [K] is the absolute

temperature, F [C mol-1] is the Faraday constant, and ν [V s-1] is the scan rate. The TOFmax

[s-1] is the �� value when �� = 1 [M]. The overpotential (η [V]) is defined by the

following eq. 3,

η = !�� "��,%��& − !�� 3�

Figure 3. A CV which exemplify the parameters of icat, ip and Ecat/2. The catalytic current increases

in the presence of 5 equivalents of acid (e.g., 2.5 mM Et3NHCl for 0.5 mM nickel(II)-based

molecular catalyst) in DMF solution. The Ecat/2 is estimated at much more positive potential than

E0P/Q.

Cur

rent

Potential

!�� !)�+�

��

��

5 eq Acid

0 eq

7

Scheme 3. The schematic flowchart for the estimation of k1 and k2 values from kglobal and Ecat/2,

obtained by electrochemical measurements. In case of EECC pathway, E and C are an electron

transfer and protonation steps, respectively.8

where !�� "��,%��& is the standard potential of the H+/H2 couple in the solvent, and !��

is the potential at the half current of catalytic current peak. Figure 3 depicts how to obtain ��, , and !�� values from a CV.

From the values of �� and !�� , the rate constant of first and second

protonation steps (i.e., �� and �) can be estimated, although the equations employed

depend on the mechanism. The flowchart for the the heterolytic EECC path is depicted in

Scheme 3. When the first PT step is the rate-determining step (RDS) (i.e., �� ≪ ��EEC′C path), “!�� ≈ !)�+� ” can be satisfied as a characteristic feature of its CV. On

��. ≈ �. �!�� ≈ !) +�� 0 �� 1 0 ��.

�.

��. ≈ ��.

!�� ≈ !) +��

�� ≪ � �� Ȃ �� is the RDSEECC′ path

�� is the RDSEEC′C path

Heterolytic EECC mechanism

E:

E:

C:

C:

O + e– P

P + e– Q

Q + H+ B

B + H+ O + H2

E0O/P

E0P/Q

k1

k2

(4)

(5)

(4-1)

(5-1)

(4-2)

(5-2)

8

the other hand, when the second PT step is rate-determining (i.e., �� Ȃ ��EECC′ path),

the !�� can be determined at more positive potential than !)�+� , because the second

term of eq. 5-2 (see Scheme 3) becomes not zero. The example of CV shown in Figure 3

can be identified as the EECC mechanism with the case of �� Ȃ � (i.e., EECC′ path),

showing large potential gap between !)�+� and !�� .

The flowchart for the case of heterolytic ECEC path is also shown in Scheme 4. As

discussed above, in this path, the first ET is coupled with the first PT (i.e., PCET). When

the catalytic currents increase with the coupling of the first PT, it suggests the possibility

of heterolytic ECEC path, which also includes the case where the second ET (at !+4�5� )

easily proceeds than the first ET (at !)�+� ). Its relationship between !�� and !)�+� is

similar to the case of EECC path. When the second ET (at !+4�5� ) is more difficult (i.e.,

more negative potential) than the first ET (at !)�+� ), the first ET (at !)�+� ) coupled with

the first PT is gradually anodically shifted with the increase of the acid concentration, and

it changes to the irreversible reduction process. Moreover, the second ET (at !+4�5� ) can

be dramatically anodically shifted with the increase of catalytic current. It should be noted

that other mechanisms have to be also considered in order to elucidate the electrocatalytic

behaviors carefully. A homolytic path or ECCE path of HER is not discussed here.

By the large contribution by Savéant and co-workers, the mechanism of

electrochemical HER by molecular catalysts has become easier to understand. The careful

analysis provides the strategy to improve the catalytic systems.

9

Scheme 4. The schematic flowchart for the estimation of k1 and k2 values from kglobal and Ecat/2,

obtained by electrochemical measurements. In the case of ECEC pathway, E and C are electron

transfer and protonation steps, resepectively.8

��. = ��.

1 0 ��.�.

!�� = !)�+� 0 �� 1 0 ��.�.

��. ≈ �.

!�� ≈ !)�+� 0�� 1 0 ��.�.

��. ≈ ��.

!�� ≈ !)�+�

�� ≪ � �� Ȃ �� is the RDSECEC′ path

�� is the RDSEC′EC path

Heterolytic ECEC mechanism

E:

C:

E:

C:

P + e– Q

Q + H+ Q'

Q' + e– B

B + H+ P + H2

E0P/Q

k1

E0Q'/B

k2

!+4�5� is easier than !)�+� !+4�5� is more difficult than !)�+�

��. = ��.

1 0 ��.�.

!�� = !+4�5� 0 �� 1 0 �.��.

“An irreversible one-electron EC wave precedes thecatalytic wave (see figure). The value of the rateconstantk1 may be derived from the positive shift ofthe peak upon addition of the substrate, whichaccompanies the passage of the wave from reversibleto irreversible”8

(6)

(7)

(8)

(9)

(6-1)

(7-1)

(6-2)

(7-2)

10

Ligand-based PCET Reduction for HER As described above, hydrogenases show high catalytic activities for HER via

consecutive metal-based PCET reductions. Researchers thus attempt to develop artificial

hydrogenase by mimicking a dinuclear structure10 or a proton-relay moiety.11 However,

it is still tough challenge to control both redox property and basicity of a metal, which are

relevant to the abilities of electron and proton transfers, in order to promote PCET for the

metal center.

In contrast, the ligand-based PCET, which is the ligand-based reduction coupled with

protonation on the ligand or metal of the coordination compound, has been paid attention

to give a flexibility for molecular designs of the catalysts for HER. The ligand-based

PCET can be rationally induced by using redox-active ligands (i.e., non-innocent ligands),

and is categorized on the basis of the previous reports (��cheme 5):

i. Reduction at the metal-ligand hybridized orbital (or ligand orbital) with a

formation of hydride species (metal-ligand-based PCET).

ii. Ligand-based PCET reduction followed by hydrogen evolution without forming

a hydride complex.

iii. Ligand-based PCET reduction followed by the formation of hydride via

intramolecular proton transfer.

Scheme 5. Ligand-based PCET leading to HER.

L(ne–)

L

+ne-,+H+

M ML(ne–)

M

+ne–,+H+

Ligand-based PCET

(i) Metal-ligand-based PCET H

H

+H+

H2

M = MetalL = Ligand

transform

nPath niii))

+H+

H2

(ii) HER from ligands

11

The mechanism of the pathway (i) has been often reported when pyridyl or π-

conjugated macrocycle (e.g. porphyrin) ligands were adopted (Scheme 6).12,13

Importantly, various metal porphyrins have catalytic activities for HER.12 For iron and

nickel porphyrins under weak acidic conditions, the reduction on porphyrin (i.e.,

M(I)P/M(I)P‒•) initiates the catalytic HER.12a-c Savéant et al. reported the catalytic

behavior of Fe(II) porphyrin for HER (Scheme 6��omplex 1).12a It evolves hydrogen from

the doubly reduced singly protonated state (i.e., Fe(II)(H)P), which forms after the second

reduction step on porphyrin affording Fe(I)P‒•. This reaction pathway can be categorized

as the EECC path, and the second protonation step was elucidated as the RDS (i.e., EECC′

path). Cao et al. succeeded to clarify the mechanism of HER by Ni(II) porphyrin (complex

2), which switches depending on the acid conditions.12c Under strong acidic conditions,

the first metal-based reduction (Ni(II/I)) initiates the formation of Ni(III) hydride. By

contrast, under weak acidic conditions, the second reduction, which was assigned as the

porphyrin-based reduction, triggers the formation of a hydride intermediate. Zhang et al.

investigated the catalytic activity and mechanism of a Ni(II) chlorin (complex 4) for HER,

which were compared with those of the Ni(II) porphyrin (complex 2).12d The Ni(II)

chlorin shows 20 times higher TOF relative to the Ni(II) porphyrin due to the difference

of hydride forming reaction, as suggested by their DFT results. For the Ni(II) porphyrin,

the hydrogenated pyrrolic nitrogen structure (i.e., ligand-based intermediate) is most

stable, in sharp contrast to the hydridonickel intermediate proposed for the Ni(II) chlorin

(Figure 4). From these results, authors emphasized that the metal-hydride-like

intermediate can more effectively promote HER via the heterolytic path in comparion

with ligand-based intermediate. Recently, Moore et al. reported the catalytic behaviors of

a copper(II) porphyrin (complex 3) and a binuclear Cu(II)2 fused porphyrin (complex

5).12e The overpotential of complex 5 was reduced by more than 500 mV compared with

complex 3, because of the large anodic shift of the reduction potential based on the

macrocyclic ligand. As the pioneering study of the molecular catalyst having pyridyl

moiety by Crabtree et al. in 1992, complex 6 was found to serve as the catalyst for HER,

which initiates by the reduction on pyridyldiimine ligand followed by the formation of

the Ni(III) hydride.13a This pioneering work also suggested the utility of non-innocent

ligands for HER.

12

Scheme 6. Molecular catalysts which proceed HER via metal-ligand-based PCET (path (i))

affording a hydride species.

Figure 4. Doubly reduced singly protonated forms of complexes 2 and 4, proposed by DFT

calculations.

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13

Catalytic systems of the HER which proceeds not via the formation of a metal

hydride species (pathway (ii)) has been recently established.14 Kato et al. reported that

iron(II) tris(o-phenylene-diamine) complex can catalyze photochemical HER (Scheme 7,

complex 7).14a,b In this system, hydrogen atom, given via photo-excited N-H dissociation,

initiates the reaction. Interestingly, Mn, Co, Ni and Zn complexes can also evolve

hydrogen by photo-irradiation, which strongly supports the hypothesis of the non-metal-

hydride mechanism. Berben et al. successfully invented the Al(III)-based molecular

catalysts for HER (complex 8),14d,g in which the mechanism was proposed to be the non-

metal-hydride pathway. They also observed the reaction intermediate by EPR, showing

the ligand-based radical coupling to two nitrogen atoms. A free unpaired electron was

suggested to delocalize on pyridyl and imino functional groups, implying the formation

of the hydrogenized ligand. Nocera and Hammes-Schiffer et al. found that mechanism of

the catalytic HER by cobalt “hangman” porphyrin (complex 9) switches from cobalt

hydride mechanism to non-cobalt-hydride mechanism depending on the acid strength.14c

Under strong acidic conditions (e.g., use of tosic acid), Co(I) can be protonated forming

Co(III) hydride. On the other hand, under weak acidic conditions (e.g., use of benzoic

acid), the HER is triggered by the second reduction of the complex 9, where its reduction

takes place over the hybridized orbital of cobalt and porphyrin. During this process, the

meso-carbon of porphyrin can be protonated followed by reacting with the pendant

carboxyl group to release hydrogen. Grapperhaus et al. reported on a Zn(II)-based

molecular catalyst promoting HER with the mechanism based on radical heterocoupling

between singly reduced singly protonated species and singly reduced doubly protonated

species (complex 10).14f It was also reported that the catalytic activity of only the ligand

is less than that of the Zn complex, which suggests that the central Zn(II) ion supports the

catalysis due to its Lewis acidity, leading to the decrease in the overpotential. For the path

(ii) shown in Scheme 5, an unsaturated amine often works not only as the proton acceptor,

but also as the hydrogen atom or hydride acceptor prior to the H2 formation.

14

Scheme 7. Catalysts for HER which proceed via no hydride forming path (path (ii)).

The case of HER via pathway (iii) is reported by only Grapperhaus et al. During the

catalytic cycle by complex 11, which has the same non-innocent ligand as complex 10,

the ligand moiety is firstly reduced by the ligand-based PCET. The subsequent reduction

occurs on the Ni center coupled with proton transfer from the ligand to the nickel center

forming a hydridonickel(III) species, which was proposed by DFT calculations (Figure

5).15

Figure 5. Intramolecular proton transfer process observed for complex 1115 triggered by the

second reduction forming a Ni(III) hydride.

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15

HER Catalyzed by Bis(dithiolene) Complexes As one of the catalysts for HER promoted by ligand-based PCET reductions, metal

bis(dithiolene) complexes have been studied for the last decades.16 Dithiolene has been

known as a redox-active ligand (i.e., non-innocent ligand) from 1960s,16a but the

motivation using this ligand was mimicking the reaction center of [NiFe] hydrogenase,

which has a nickel tetrathiolate structure.5

One of the early studies on metal dithiolene complexes as the catalysts for HER was

done by Kisch et al.,17a in which they reported on the photocatalytic activities of several

metal bis or tris(dithiolene) complexes for HER (Scheme 8a). Also in 1980, Vlček et al.

reported the electrocatalytic behaviors of bis(mnt) complexes of Ni, Co, and Rh for HER

(mnt = maleonitriledithiolate) (Scheme 8b)17b where the mechanism of the HER by

[M(mnt)2]2– was firstly discussed on the basis of the electrochemical and spectroscopic

measurements (see below).

Scheme 8. Metal dithiolene complexes studied for their catalytic abilities for (a) photochemical

HER in 1980 by Kisch et al.,17a and (b) electrochemical HER in 1980 by Vlček et al.17b

Sakai et al. reported that [FeIII (mnt)2]22– (Scheme 9) has a catalytic activity for HER

in an aqueous acidic buffer solution in 2009.18a This is the first report realizing the HER

by metal dithiolenes in fully aqueous media. It was found that the HER by [FeIII (mnt)2]22–

proceeds from the acid as the proton source (i.e., acetic acid in this work), not from H2O.

They also succeeded to improve the system of HER, in which [FeIII (dcpdt)2]22– (dcpdt =

5,6-dicyanopyrazine-2,3-dithiolate��9) catalyst does not require any acids as the

proton source to promote electrochemical HER in water.

(a) (b)

16

Scheme 9. Molecular structures of [Fe(mnt)2]22– (left) and [Fe(dcpdt)2]2

2– (Right).

Following these pioneering works, the catalytic activities of Fe,18 Co,19 Ni,20 Mo,21

and W22 bis(dithiolene) complexes for HER were investigated, and some reports also

clasified the mechanisms by electrochemical and computational studies (see next section).

Bis(dithiolene) complexes have been studied not only as the homogeneous catalysts,

but also as catalysts based on 1D or 2D metal-organic frameworks strongly promoted by

Marinescu et al.23,24 The catalytic activity of bis(dithiolene) complexes for CO2 reduction

reaction, which evolve formic acid as the main product via a hydride pathway, was

recently investigated by Fontecave et al.25 Catalysis of bis(dirthiolene) complexes are

now open for various reduction reactions.

Mechanism of HER Catalyzed by Bis(dithiolene) Complexes As discussed above, it was already found in 1980 that metal dithiolene complexes

have catalytic activities for HER. Interestingly, the mechanism was also discussed at that

time. Specifically, Vlček et al. discussed the mechanism of HER by a bis(dithiolene)metal

complex ([Rh(mnt)2]2– in particular).17b They succeeded to observe the formation of

singly reduced singly-protonated species after the first reduction of [Rh(mnt)2]2–, which

was assigned as the hydridorhodium(III) species. It was also suggested that the hydride

species is further reduced, leading to the hydrogen evolution.

In 2012, Hammes-Schiffer and Solis discussed the mechanism of [Co(bdt)2]– (bdt =

benzene-1,2-dithiolate�� 6a) series and [Co(mnt)2]22– (Figure 6b) by using DFT

calculations.19c Catalytic activities of these compounds were previously reported by

Eisenberg et al.,19a,b where it was briefly suggested that the mechanisms adopt ECEC path.

Hammes-Schiffer and Solis carefully assigned each electrochemical behavior by DFT,

and found that the first ligand-based reductions of the complexes couple with one or two

proton transfer to thiolate or metal. The initial ligand-based reduction of [Co(bdt)2]– is

17

coupled with two proton trasfer forming the Co(III) hydride with one of the tholates

protonated, followed by the reduction of the singly reduced doubly protonated species in

the presence of trifluoroacetic acid (Figure 6a). The electrocatalytic HER by [Co(bdt)2]–

series proceeds via the E(CC)EC path. On the other hand, one-electron-reduced product

of [Co(mnt)2]– having stronger electron-withdrawing groups can be singly protonated at

a sulfur donor, subsequently undergoing the reduction of the cobalt center followed by

the intramolecular proton transfer to form the Co(III) hydride species (Figure 6b). This

path can be clasified as the ECEC mechanism. The impact of this work is that (i) the

thiolate can work as the proton relay, (ii) the reduction potential shifts anodically by using

the dithiolene ligands having electron-withdrawing groups, which is relevant to reduce

the overpotential, and (iii) this ligand also reduces the basicity of the thiolate suppressing

its PCET. These observations suggested that the activity and mechanism of HER by

bis(dithiolene) complexes are strongly affected by the electronic properties of dithiolene

ligands.

Figure 6. The proposed mechanism of HER catalyzed by metal (M = Co, Ni) bis(dithiolene)

complexes having electron-donating groups (a) or electron-withdrawing groups (b).

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(a) Electron-donating groups

=

(b) Electron-withdrawing groups

=

18

Mitsopoulou et al. also succeeded to cralify the mechanisms of [Ni(bdt)2]– and

[Ni(mnt)2]– family for HER.20d,e Their DFT results showed similar catalytic pathways to

the cobalt series, but in the case of nickel complexes, the formation of hydridonickel

intermediate is relatively unfavorable than the protonation on the sulfur. The mechanisms

of both [Ni(bdt)2] – and [Ni(mnt)2]– do not include the steps affording hydride species.

These DFT studies suggested that the protonation site for the ligand-based PCET

reduction is on the thiolate (i.e., on the ligand). However, the possibility of the formation

of a metal hydride species as the following process during the catalytic HER is not still

discussed in detail. Cordones et al. predicted the formation of a metal hydride species by

the experimental observations during the electrocatalytic HER by Ni(adt)2 (adt = 2-

aminobenzenethiolate��. Ni(adt)2 is not a bis(dithiolene) complex, but has similar

property because of its ligand non-innocence.25 The X-ray absorption spectroscopy

(XAS) and electronic structure calculations revealed that the LUMO of Ni(adt-H)2, in

which one more electron reduced Ni(adt-H)2 (i.e., [Ni(adt-H)2]-) is considered as a key

intermediate for HER, is mainly localized on the metal and bonded-S/N atoms (i.e., Ni-

S/N). This result suggested that the reduction of this LUMO leads to formation of the Ni-

H intermediate followed by H2 elimination (Figure 7). Most of the intermediates have not

been observed or isolated yet, but this observation proposed the necessity of forming a

metal hydride species for the HER by bis(dithiolene) molecular catalysts.

Figure 7. Proposed mechanism of catalytic HER by Ni(adt-H)2

Ni(adt-H)2 [Ni(adt-H)2]- [Ni(-H)(adt)(adt-H)]-

19

Survey of This Thesis As discussed above, several examples of metal bis(dithiolene) complexes

catalyzing photochemical and electrochemical HER have been developed. Some

computational studies revealed that the HER proceeds via several types of the ligand-

based PCET. The molecular system of metal bis(dithiolene) catalysts is one of the

most studied examples as the catalyst for HER, promoted by the ligand-based PCET

reductions. However, as briefly discussed in the previous section, the dithiolene

ligands as the ligand-based PCET acceptor have several unsolved issues in order to

develop the highly active molecular systems for HER. Firstly, the protonation on

thiolate is not favored due to its low basicity (i.e., low pKa). Some of the previous

studies evaluating the catalytic activities of bis(dithiolene) molecular catalysts

employed strong acids to promote the ligand-based PCET reduction over dithiolene

ligands. There has been almost no examples studying the catalysis under weak acidic

conditions. Secondly, the basicity of thiolates is strongly reduced by the presence of

electron-withdrawing substituents, where the positive shift of the reduction potential

suppresses the increase in pKa of thiolates leading to less promotion of PCET. In

order to realize the catalytic activity comparable to hydrogenases, development of

the effective catalytic system under neutral conditions is highly required. In addtiton,

promotion of the ligand-besed PCET reduction over the dithiolene ligands is also a

crucial target to achive the artificial hydrogenase.

In this context, the main focus of this thesis is on the development of the

effective ligand-based PCET systems for electrochemical HER based on metal

dithiolene complexes. In addition, the mechanisms of catalytic HER are analyzed by

electrochemical and theoretical studies. Knowledge obtained by these studies is

expected to be applied for the design of new catalysts having non-innocent ligands.

20

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23

Chapter 1: A Nickel Dithiolate Water Reduction

Catalyst Providing Ligand-based Proton-coupled

Electron Transfer Pathways

Introduction Artificial photosynthesis driving water splitting (2H2O � 2H2 + O2)1-4 is a key to

establish technologies enabling solar energy conversion into hydrogen energy, the

simplest form of clean energy affording water upon either combustion or oxidation in fuel

cells. In order to fabricate systems enabling higher solar energy conversion efficiency,

development of fast, robust catalysts operating with minimized overpotentials (OPs) is a

crucial target. Moreover, the use of earth abundant metal ions together with the avoidance

of volatile organic compounds would enable the widespread, practical applications.

Scheme 1. The active centers of [NiFe] and [FeFe] hydrogenases.

For the hydrogen evolution reaction (HER), nature invented efficient molecular

systems, [NiFe] and [FeFe] hydrogenases (see Scheme 1).5 They catalyze HER with

a nearly ��-

1 and 6000-21000 s-1 for the [NiFe]5a and [FeFe]5a,5c hydrogenases, respectively. In

general, catalysts showing a high TOF with a low OP is considered as an efficient

catalyst.2g The high TOFs observed for the [FeFe] hydrogenase have been considered

to largely rely on a pendant amine donor located in close proximity to the active iron

center where hydride formation as well as coupling of a hydride and a proton rapidly

[NiFe] hydrogenase [FeFe] hydrogenase

24

takes place (M + H+ + 2e- � M(H)- and M(H)- + H+ � M + H2, where M is an active

iron center) (see Scheme 1).5d,6 Fast proton delivery to the active center is

substantially enhanced by the proton relay using the pendant amine donor. Some

successful examples of Ni-based molecular catalysts bearing such proton relay sites

(see Scheme 2) have been developed by DuBois and co-workers.7 Moreover, an

increasing number of non-precious metal H2-evolving catalysts have been reported.2,3

Nevertheless, compounds capable of catalyzing HER in aqueous media free of

organic solvents are still rare.2c,4,8-10 For a water-soluble Co-NHC H2-evolving

catalyst (NHC = N-heterocyclic carbene�� Scheme 2), Sakai et al. recently

highlighted its relatively low OP for HER (onset-��

(proton-��

Co(II) + H+ + e‒ � Co(III)(H) (presumably, metal-centered PCET) via concerted

electron filling and protonation at the dz2 orbital.10a This path can avoid preliminary

metal-centered reduction which often requires a relatively high OP. This is in sharp

contrast with the generally reported Co(III)(H) formation triggered by simple

protonation of a low valent d8-��‒ � Co(I) and Co(I) + H+ �

Co(III)(H).2b,11

In contrast with the protonation at the filled dz2 orbital of the d8-Co(I) species,

simple protonation over the Pt(II) or Ni(II) d8 ions is unfavorable. Nevertheless,

Yamauchi and Sakai recently demonstrated that the HER catalyzed by [PtCl(tctpy)]2‒

(tctpy = terpy-��Scheme 2) can be triggered by the ligand-centered

reduction accompanied by protonation at one or two carboxylates on the terpyridine

ligand (i.e., ligand-centered PCET).10b In this system, a hydridoplatinum(III)

intermediate is considered to be given via reduction by the electron stored over the

tctpy ligand bound to the Pt(II) ion (i.e., [PtIICl(tctpyHn‒•)](3-n)‒ + H+ →

PtIII (H)Cl(tctpyHn)](2-n)‒).4c

One of the interests over the last decade has also concentrated on the HER

catalyzed by transition-metal dithiolenes having an MS4 core to develop artificial

hydrogenase mimics by employing the sulfur donors from dithiolate ligands. The

pioneering study on the HER by this family was reported by Kisch and co-workers

in 1980, in which the Ni, Pd, Pt, Fe, Co, Mo, and W dithiolenes were examined for

25

photocatalytic HER.12a Since then, several researchers also investigated the HER by

the ditholene complexes of Fe,9a-e,12b,c Co,12d,e Ni,12f,g,h Mo,12i Rh,12j and W.12k These

involve the initial study from Sakai group on a dinuclear iron dithiolene catalyst

[FeIII (mnt)2]22‒ (mnt = maleonitriledithiolate).9a It was also realized that

[FeIII (dcpdt)2]22- (dcpdt = 5,6-dicyanopyrazine-2,3-dithiolate) exhibits an improved

catalytic performance in that it does not require the presence of any acid source

except for water, to catalyze HER in alkaline aqueous media (pH = 11),9b-e although

acetic acid is required for the mnt derivative.9a These results suggested that the

pyrazine donors can abstract protons from water molecules during the catalytic cycles.

Here a new Ni-based molecular catalyst ([Ni(dcpdt)2]2‒) for electrochemical

HER, having a NiS4 core (Scheme 2), is reported. [Ni(dcpdt)2]2‒ catalyzes HER with

relatively low OPs, likely due to its unique ligand-based reduction accompanied by

protonation of the pyrazine donors (i.e., ligand-based PCETs) leading to the nickel-

hydride intermediates without forming low-valent Ni(I) or Ni(0) species. Such

ligand-based reduction processes have never been discussed in the previous reports.

Scheme 2. Molecular catalysts for HER.

��

� �

�

�

��

��

��

��

��

��

��

��

�

�

�

��

��

�

�

�

� ��

��

��

��

��

��

��

�

�

��

��

��

��

�

�

��

��

��

��

26

Experimental Section

Materials NiCl2•6H2O and Na2S•9H2O were purchased from Kanto Kagaku. 2,3-dichloro-5,6-

dicyanopyrazine was purchased from Tokyo Chemical Industry. All solvents and reagents

were of the highest qualities available and were used as received without further

purification.

Synthesis of Na2[Ni(dcpdt) 2]•4H2O Water-soluble sodium salt of [Ni(dcpdt)2]2- (Na2[Ni(dcpdt)2�� dcpdt = 5,6-

dicyanopyrazine-2,3-dithiolate) was prepared according to the method reported for

[NBu4]2[Ni(dcpdt)2],13 with minor modifications as follows. To a stirring suspension of

sodium sulfide nonahydrate (1.99 g, 8.4 mmol) in acetone (200 mL) was added a solution

of 2,3-dichloro-5,6-dicyanopyrazine (0.537 g, 2.7 mmol) in acetone (50 mL). After

stirring for 1 h at room temperature, the resulting red orange mixture was filtered for the

removal of unreacted solids. To the filtrate was added a solution of nickel chloride

hexahydrate (0.315 g, 1.34 mmol) in methanol (50 mL). After stirring for 5 min, the

resulting black mixture was filtered off, and diethyl ether was slowly added to the filtrate.

The black purple product deposited was collected by filtration. The solid was dissolved

in hot water (200 mL) and the solution was filtered for the removal of insoluble materials.

The filtrate was evaporated to dryness. The same procedure was performed by dissolving

the solid in acetone (80 mL) to afford the final product as a black purple powder (yield:

0.318 g, 0.62 mmol, 23%). Anal. Calcd for C12N8NiS4•4H2O (f.w. 515.18): C, 25��

��13C NMR (D2O, 600 MHz): 116.27,

121.81, 177.76.

Measurements 13C NMR spectrum was acquired on a JEOL JNM-ESA 600 spectrometer. Cyclic

and square wave voltammetry and bulk electrolysis were performed on a BAS ALS

602DKM electrochemical analyzer. For these experiments, a glassy carbon (GC) or

indium tin oxide (ITO) working electrode, a platinum wire counter electrode, and a

��where the

27

ITO electrode was purchased from BAS (No. 010887). The bulk electrolysis was carried

out by an H-type cell (VB-9) purchased from EC Frontier, using a GC rod working

electrode (5 mm Φ, The Nilaco Corporation), a platinum mesh counter electrode, and a

SCE. The working compartment was separated from the counter compartment using a

cation exchange membrane (SelemionTM CMD, AGC Engineering). The time-course of

H2 evolution during the bulk electrolysis was monitored using the automated system

developed in Sakai group. These experiments adopted the continuous Ar-flow method (10

mL min-1) with the vent introduced into the auto sampler for the gas chromatographic

analysis, as described elsewhere.14 The pH measurements were performed using a DKK-

TOA HM-25R pH meter. Energy dispersive X-ray fluorescence (EDX) spectrum was

recorded using a Shimadzu EDX-720 spectrometer with a Rh target.

DFT calculations In order to better understand the structural and spin-state candidates, density

functional theory (DFT) calculations were performed using the Gaussian 09 package of

programs.15 The structures were fully optimized using the B3P86 density functional16,17

with the effect of solvation in water taken into consideration using the conductor-like

polarizable continuum model (C-PCM) method.18,19 The 6-311+G(d,p) basis set was

applied to all atoms. The use of B3P86 functional was reported to show good consistency

with theoretical and experimental results for the 1st row transition metal complexes,20

which is also continued in the extensive studies attempting to clarify the mechanism of

hydrogen evolution reaction (HER) by the present system. The details will be separately

reported in Chapter 2 of this thesis.

28

Results and Discussion Figure 1A shows cyclic voltammograms (CVs) recorded for aqueous solutions of

[Ni(dcpdt)2]2‒ at three different pH conditions (pH = 4.0, 5.0, and 6.0), all showing

distinct flow of cathodic current based on HER catalyzed by this molecular catalyst

(similar behaviors are seen in the CVs at pH = 8.0 and 9.0� Figure 2).

Figure 1. A) CVs for aqueous acetate buffer solutions (pH = 4-��2[Ni(dcpdt)2]•4H2O

(0.5 mM) containing NaCl (0.1 M) at room temperature under Ar atmosphere, recorded at a sweep

rate of 100 mV/s. The working, counter, and reference electrodes were a glassy carbon (GC) disk,

a Pt wire, and a saturated calomel electrode (SCE), respectively. B) SWVs of

Na2[Ni(dcpdt)2]•4H2O (0.5 mM) in the same condition as in Figure 1A.

-1.2 -1 -0.8 -0.6 -0.4 -0.2 0

a) [Ni(dcpdt)2]2- (pH = 4)

b) [Ni(dcpdt)2]2- (pH = 5)

c) [Ni(dcpdt)2]2- (pH = 6)

Potential / V vs. SCE

e) Blank (pH = 5)

b) [Ni(dcpdt)2]2- (pH = 5)

d) Blank (pH = 4)

a) [Ni(dcpdt)2]2- (pH = 4)

f) Blank (pH = 6)

c) [Ni(dcpdt)2]2- (pH = 6)

Cur

rent

/ µA

A)

B)

5 µA

5 µA

29

��

��

��

��

��‒��

��η(��η for HER,��

��

��

��

��

��

30

[a] These values were estimated from the CVs and linear sweep voltammograms of Na2[Ni(dcpdt)2]•4H2O in 0.1 M

aqueous acetate buffer (pH = 4-6) and borate buffer (pH = 8-9) solutions containing NaCl (0.1 M) (see also Figure 3).

From Ecat/2 values determined for the catalytic currents for electrochemical HER in

Figure 3, the η(Ecat/2) values for HER by [Ni(dcpdt)2]2‒ are measured to be 330-400 mV

as summarized in Table 1. It is considered that this is a rare example of molecular catalyst

exhibiting such low η(Ecat/2) values in aqueous media free of organic solvents. Important

examples showing low overpotentials for HER involve bis(diphosphine)nickel

complexes (η(Ecat/2) = 250-750 mV8b,c,g) and a self-assembled cobalt complex (η(Ecat/2)

not reported).8d

As shown in Figure 4, electrochemical HER catalyzed by [Ni(dcpdt)2]2‒ was also

�� by the controlled potential

electrolysis at different applied OPs in the range 0.31-0.41 V. The amount of H2 evolved

dramatically increases by the presence of [Ni(dcpdt)2]2‒ (Figure 4b,d). As expected, the

rate of HER changes dramatically upon changing the applied overpotential. When HER

is driven by 0.36 and 0.41 V in overpotential, [Ni(dcpdt)2]2‒ exhibits high activity for

HER with the TON (24 h) reaching 16000 and 20000, respectively (Table 2), indicating

its excellent durability in a long-term use in electrochemical HER. The Faradaic

efficiency has been estimated as 92-100% (Table 2), also showing its distinct property as

an excellent catalyst for HER.

Table 1. pH dependence of η(Ecat/2) values for HER catalyzed by [Ni(dcpdt)]2‒.[a]

pH 4.0 5.0 6.0 8.0 9.0

Standard electrode potential for HER /V vs. SCE -0.48 -0.54 -0.60 -0.71 -0.77

OP(Ecat/2) / V

(Ecat/2 / V vs. SCE)

0.33

(-0.81)

0.37

(-0.91)

0.40

(-1.0) N.D. N.D.

31

Figure 3. A,B,C) Linear sweep voltammograms (LSVs) for aqueous acetate buffer solutions (pH

= 4-�� 2[Ni(dcpdt)2]•4H2O (0.5 mM) containing NaCl (0.1 M). η(Ecat/2)’s are

determined as illustrated in Figures. D) LSVs of Na2[Ni(dcpdt)2]•4H2O (0.5 mM) recorded at pH

= 8,9 (0.1 M aqueous borate buffer solutions containing 0.1 M NaCl), showing no peak currents

observed that do not allow us to determine Ecat/2. All the LSVs are recorded at a sweep rate of

250 mV/s, under Ar atmosphere at room temperature.

-5

5

15

25

35

45

55-1.4 -1.3 -1.2 -1.1 -1 -0.9 -0.8 -0.7 -0.6

Cur

rent

(µA

)

Potential (V vs. SCE)

pH = 8.0

pH = 9.0

D) pH = 8.0, 9.0

-5

5

15

25

35

45

55

65-1.2 -1.1 -1 -0.9 -0.8 -0.7 -0.6 -0.5 -0.4

Cur

rent

(µA

)


icat

icat/2

η(Ecat/2)

A) pH = 4.0

0.33 VEcat/2

E2H+/H2

-5

5

15

25

35

45-1.2 -1.1 -1 -0.9 -0.8 -0.7 -0.6 -0.5 -0.4

Cur

rent

(µA

)


B) pH = 5.0

icat

icat/2

Ecat/2

η(Ecat/2)0.37 V

E2H+/H2

-5

0

5

10

15

20

25-1.3 -1.2 -1.1 -1 -0.9 -0.8 -0.7 -0.6 -0.5

Cur

rent

(µA

)


C) pH = 6.0

icat

icat/2

Ecat/2

E2H+/H2

η(Ecat/2)0.40 V

32

Figure 4. Electrochemical H2 evolution catalyzed by Na2[Ni(dcpdt)2]•4H2O (1 µM) during the

controlled potential electrolysis at a) -0.85 V, b) -0.90 V, and c) -0.95 V vs. SCE in aqueous

��

The working, counter, and reference electrodes were a GC rod, a Pt wire, and a SCE, respectively.

0.0

1.0

2.0

3.0

4.0

5.0

6.0

0 5 10 15 20

a) -0.85 V vs. SCE (OP = 0.31 V)b) -0.90 V vs. SCE (OP = 0.36 V)c) -0.95 V vs. SCE (OP = 0.41 V)d) Blank at -0.90 V vs. SCE (OP = 0.36 V)

H2 e

volv

ed /

mL

Electrolysis time / h

33

Table 2. TONs and Faradaic efficiencies for the bulk-electrolysis of a 0.1 M aqueous acetate

buffer solution (pH = 5.0) of [Ni(dcpdt)2]2‒ (1 µM). See Figure 4 for details.

Electrolysis potential / V vs. SCE

(Applied overpotential/ V)

-0.85

(0.31)

-0.90

(0.36)

-0.95

(0.41)

TON of [Ni(dcpdt)2]2‒ (24 h) 3300 16000 20000

Faradaic efficiency 92% 95% ≈100%

The pH-dependent redox behaviors of [Ni(dcpdt)2]2‒ were extensively studied in

order to gain insights into the mechanism of HER. Figure 5 shows the Pourbaix diagram

developed based on the square wave voltammograms (SWVs) observed for its aqueous

buffer solutions at pH = 3.2-6.4 (see Figures 1B, 6, 7 and Tables 3, 4). It is noted that the

data at pH < 3.2 are not observable due to deposition of protonated species, while those

at pH > 6.4 are reported elsewhere.12l A pH-independent oxidation process of

[Ni(dcpdt)2]2‒ was observed at pH > 5.0. The first oxidation process for such Ni

dithiolates have been shown to proceed as ligand-based 1-electron oxidation.22 From the

slope of the pH-��-55 mV/pH decade) and pH-independent

lines for this 1-electron process, the proton dissociation constant for eq. 2 can be estimated

(pKa = 5.0).

[Ni II(dcpdt)(dcpdtH)]‒ ⇄ [Ni II(dcpdt)2]2‒ + H+ pKa = 5.0 (2)

34

Figure 5. Plot of the first reduction and oxidation potential of [Ni(dcpdt)2]2‒ as a function of pH

(Pourbaix diagram), where potentials were determined by observing SWVs of the complex in

aqueous media at various pH conditions. See Figures 1B, 6, 7 and Tables 3, 4 for details.

-1.6

-1.2

-0.8

-0.4

0

0.4

0.8

3 4 5 6

Pot

entia

l / V

vs.

SC

E

pH

[NiII(dcpdt)2]2‒

E = +0.39

ET (e-)

[NiII(dcpdt)(dcpdt+•)]‒

[NiII(dcpdt)(dcpdtH)]‒

pKa = 5.0

[NiII(dcpdt)(dcpdtH2)]‒[NiII(dcpdtH)2]‒

or

35

��

��

��

36

��

��

��

��

��

37

pH Oxidation potential

/ V vs. SCE

3.21 0.498

3.6 0.470

4.09 0.420

4.3 0.416

4.64 0.412

4.8 0.404

4.99 0.400

5.2 0.392

5.89 0.392

Table 3. Oxidation potentials of Na2[Ni(dcpdt)2]•4H2O evaluated from the SWVs shown in

Figure 6. These values are employed to draw the Pourbaix diagram depicted in Figure 5.

38

Table 4. Reduction potentials of Na2[Ni(dcpdt)2]•4H2O evaluated from the SWVs shown in

Figure 7. These values are employed to draw the Pourbaix diagram depicted in Figure 5.

pH Reduction potential

/ V vs. SCE

3.60 -0.832

4.09 -0.860

4.30 -0.876

4.46 -0.880

4.64 -0.900

4.80 -0.900

4.99 -0.920

5.20 -0.950

5.32 -0.936

5.41 -0.992

5.61 -0.996

5.78 -1.024

5.89 -1.071

6.06 -1.072

6.32 -1.104

6.35 -1.072

6.40 -1.088

39

On the other hand, the first reduction process can be classified into two pH domains

exhibiting different types of reduction processes. At pH < 5.0, [NiII(dcpdt)(dcpdtH)]‒

undergoes 1-electron reduction via PCET coupled with one proton and one electron

transfer (abbreviated as PCET(H+,e‒), eq. 3), since the slope of the first reduction is

determined as -62 mV/pH decade. Meanwhile, the behavior in the pH = 5.0-6.4 domain

indicates that PCET(2H+,e‒) (eq. 4) proceeds for [NiII(dcpdt)2]2‒ from the slope of -127

mV/pH decade.

[Ni II(dcpdt)(dcpdtH)]‒ + H+ + e‒ ⇄

[Ni II(dcpdt)(dcpdtH2)]‒ or [NiII(dcpdtH)2]‒ (3.6 < pH < 5.0) (3)

[Ni II(dcpdt)2]2‒ + 2H+ + e‒ ⇄

[Ni II(dcpdt)(dcpdtH2)]‒ or [NiII(dcpdtH)2]‒ (5.0 < pH < 6.4) (4)

As discussed above, the catalyst-based reduction peaks, unobservable in CV scans

(Figure 1A), are indeed observable in SWV scans (Figure 1B), consistent with the

conclusion that the electrochemical HER by [Ni(dcpdt)2]2‒ is triggered by the 1-electron

reduction of the catalyst.

40

Several additional experiments were carried out to ascertain that the observed

catalytic HER is not derived from catalysis by any undesirable side products23 that might

be formed during the catalysis. The standard “rinse test”, conducted using a GC working

electrode, revealed that the electrode used in the 100-cycles of cathodic sweep for HER

exhibited only a minor catalytic effect for HER when the electrolysis solution was

replaced with a solution free of [Ni(dcpdt)2]2‒ (see Figure 8f). Moreover, the careful

examinations revealed that this nickel catalyst is adsorbed over the GC electrode surface

by merely soaking the electrode in a catalyst solution (Figure 8g). Similarly, it was also

confirmed the catalyst adsorption over the GC rod electrode after the bulk-electrolysis

(Figure 4b). The molar ratio of S and Ni involved in the materials adsorbed over the GC

rod was determined as 3.8 by EDX (Energy Dispersive X-ray fluorescence) spectroscopy,

Figure 8. ��

pH = 5) of Na2[Ni(dcpdt)2]•4H2O (0.5 mM) in the presence of NaCl (0.1 M), where 1st, 5th, 25th,

50th, and 100th cycles are only shown. The f-labeled CV shows the result of rinse test after the

e-labeled scan, recorded after replacing the electrolysis solution with the same buffer solution

free of the catalyst. The g-labeled CV was recorded by the same GC electrode which was

preliminary soaked in a catalyst solution, where the measurement was carried out using the same

buffer solution free of the catalyst. The result reveals that the catalyst has a tendency to be

adsorbed over the GC electrode without any potential sweeps. All the CVs are recorded at a sweep

rate of 100 mV/s, under Ar atmosphere at room temperature.

-2.5

0

2.5

5

7.5

10

12.5-1 -0.8 -0.6 -0.4 -0.2 0

Cur

rent

(µA

)


blank

a) 1st cycle

b) 5th cycle

c) 25th cycle

d) 50th cycle

e) 100th cycle

f) rinse test

g) soaked

5

7

9

11

-0.98 -0.94 -0.9

1st cycle

100th

blank

a) 1st cycle

b) 5th cycle

c) 25th cycle

d) 50th cycle

e) 100th cycle

f) rinse test

g) dip-coated

41

revealing that the NiS4 core is preserved in the materials adsorbed. The specific affinity

of the catalyst over the GC surfaces was evidenced by the fact that such adsorption can

be completely suppressed when an indium tin oxide (ITO) electrode was employed as a

working electrode (Figure 9). By use of the ITO electrodes, neither simple soaking nor

100 cycles of scanning for HER did not result in adsorption of neither metallic nor

molecular materials which could not be removed in the subsequent rinse test (Figure 9f,g).

These results reveal that degradation of [Ni(dcpdt)2]2‒ is negligible and the catalytic HER

by [Ni(dcpdt)2]2‒ undergoes as homogeneous catalysis.

Figure 9. CVs with use of an ITO working el��

pH = 5) of Na2[Ni(dcpdt)2]•4H2O (0.5 mM) in the presence of NaCl (0.1 M), where 1st, 5th, 25th,

50th, and 100th cycles are only shown. The f-labeled CV shows the result of rinse test after the

e-labeled scan, recorded after replacing the electrolysis solution with the same buffer solution

free of the catalyst. The g-labeled CV was recorded by the same ITO electrode which was

preliminary soaked in a catalyst solution, where the measurement was carried out using the same

buffer solution free of the catalyst. The result reveals no materials are adsorbed over the ITO

electrode even with 100 cycles of CV scans or simple soaking of the ITO electrode to the catalyst

solution. All the CVs are recorded at a sweep rate of 100 mV/s, under Ar atmosphere at room

temperature.

-5

5

15

25

35-1 -0.8 -0.6 -0.4 -0.2 0

Cur

rent

(µA

)


15

20

25

30

35-0.98 -0.94 -0.9

1st cycle

100th

blank

a) 1st cycle

b) 5th cycle

c) 25th cycle

d) 50th cycle

e) 100th cycle

f) rinse test

g) dip-coated

42

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�� ‒� ��

��

��

��

��

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��

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��

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44

indicates that the first PCET reduction of [NiII(dcpdt)2]2‒ undergoes as the ligand-based

PCET, and it triggers the HER. The pyrazine moiety promotes the PCET process, also

leading to the decrease in overpotential of HER. This is a quite rare example of the

molecular catalysis for HER, and extended studies including elucidation of the whole

mechanism by DFT will be discussed in the next chapter.

45

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48

Chapter 2: Consecutive Ligand-based PCET Processes

Affording a Doubly Reduced Nickel Pyrazinedithiolate

which Transforms into a Metal Hydride Required to

Evolve H2

Introduction Towards the development of a sustainable energy society, considerable attention

has been paid to the studies on solar water splitting reactions (2H2O → 2H2 + O2).1-3

In order to fabricate systems enabling higher solar energy conversion efficiency,

development of fast, robust catalysts operating with a minimized overpotential is one

of the crucial targets.4 The study here is focused on the water reduction side of

catalysis using transition metal molecular systems.

Until now, various complexes of cobalt,1a,2b,2e,4a,5 nickel2b,2d,2e,4a,5c,6 and

platinum7 have been studied in detail for their activity and mechanism in hydrogen

evolution reaction (HER). As depicted in Scheme 1A, the HER’s by molecular

catalysts are considered to proceed via four types of paths leading to a hydride

intermediate. Many reports so far employed the type-1 route in which reduction of

the metal-based orbital is followed by the protonation to give a metal hydride

intermediate.1a,2e,5a This route is often denoted as an ET-PT (Electron Transfer Proton

Transfer) path.8 Some reports also adopted the type-3 route in which reduction of the

metal-ligand hybrid orbital is a key to afford a metal hydride.6b,6d,7b,7c It is considered

that reactions may also fall into type 2 in which neither PT-ET nor ET-PT path is

allowed but only the proton-coupled electron transfer (PCET)8,9 path is allowed to

give a metal hydride, which may be called ‘Metal-centered PCET’ path.5b Finally, it

has been recently found that [NiII(dcpdt)2]2- (depicted in Scheme 1B) is the first

example of catalyst which adopts the type-4 route in which the ligand-centered PCET

reduction is followed by the intramolecular relocation of proton and electrons to give

a metal hydride.10a This finding shed a new light in the field of molecular catalysts

49

��

��

��

��

��

��

��

��

��

��

��

��

��

��

��

��

50


Computational Method Density functional theory (DFT) calculations were performed using Gaussian

09/16 packages11 to understand the structural and spin-state candidates. The

structures were fully optimized using the B3P86 density functuonal12,13 with the

effect of solvation in water taken into consideration using the conductor-like

polarizable continuum model (C-PCM) method.14,15 The 6-311+G(2d,p) basis set was

applied to all atoms. The use of B3P86/6-311+G(2d,p) level of DFT was reported to

show good consistency with theoretical and experimental results for the 1st row

transition metal complexes,16 which were further confirmed by the calculations on

model systems (data not shown). The redox potentials and pKa values are calculated

as described in Scheme 2. The structure of the transition state was determined using

the TS and QST3 methods, followed by performing the intrinsic reaction coordinate

(IRC) calculations.

51

Scheme 2. Isodesminc reaction methods16,17 for calculating redox potentials and pKa values based

on the experimentally determined values for the oxidation potential of [NiII(dcpdt)2]2‒ (+0.39 V

��5,6-dicyanopyrazine-2,3-dithiolate) and pKa for [Ni II(dcpdt)(dcpdtH)]‒ (pKa =

5.0), respectively. F is Faraday constant, R is the gas constant, and T is temperature (298.15 K).

(1) Determination of the redox potential for Ox + e– → Red

Ox + e– → Red ∆G0 = –FE0

[Ni II(dcpdt)2]2‒ → [Ni II(dcpdt)(dcpdt+•)]‒ + e–

∆G0ref = –FE0

ref E0ref = 0.39 V vs. SCE

Ox + [NiII(dcpdt)2]2‒ → Red + [NiII(dcpdt)(dcpdt+•)]‒

∆G0r = G(Red)solv + G([Ni II(dcpdt)(dcpdt+•)]‒)solv – G(Ox)solv – G([Ni II(dcpdt)2]2‒)solv

∆G0r = –∆G0

ref + ∆G0 = FE0ref – FE0

E0 = –∆G0r /F + E0

ref

(2) Determination of pKa value for A + H+ → AH+

A + H+ → AH+ ∆G0 = RTln(10)pKa

[Ni II(dcpdt)2]2‒ + H+ → [Ni(dcpdt)(dcpdtH)]‒

∆G0ref = RTln(10)pKa,ref pKa,ref = 5.0

A + [Ni II(dcpdt)(dcpdtH)]‒ → AH+ + [NiII(dcpdt)2]2‒

∆G0r = G(AH+)solv + G([Ni II(dcpdt)2]2‒)solv – G(A)solv – G([Ni II(dcpdt)(dcpdtH)]‒)solv

∆G0r = RTln(10)pKa – RTln(10)pKa,ref

pKa = ∆G0r/RTln(10) + pKa,ref

52

Measurements Linear sweep voltammetry (LSV) was performed on a BAS ALS 602DKM

electrochemical analyzer and a BAS RRDE-3A Rotating Ring Disk Electrode

Apparatus. For the experiments for the aqueous solutions, a glassy carbon working

electrode (5 mmφ), a platinum wire counter electrode and a saturated calomel

��vs. NHE) were employed, where NaCl (0.1 M)

was used as a supporting electrolyte. For the experiments for the organic solutions, a

glassy carbon working electrode (5 mmφ), a platinum wire counter electrode and an

Ag/Ag+ reference electrode (0.249 V vs SCE) were employed, where TBAPF6

(tetra(n-��

electrolyte and all potentials reported are given relative to the Fc/Fc+ couple (Fc/Fc+

= 0.155 vs SCE). The bulk electrolysis was carried out by an H-type cell (VB-9)

purchased from EC Frontier, using a GC rod working electrode (5 mm φ, The Nilaco

Corporation), a platinum plate counter electrode, and an SCE. The working

compartment was separated from the counter compartment using a cation exchange

membrane (SelemionTM CMD, AGC Engineering). The time-course of H2 evolution

during the bulk electrolysis was monitored using the automated system developed in

Sakai group. These experiments adopted the continuous Ar-flow method (10 mLmin‒

1) with the vent introduced into the auto sampler for the gas chromatographic analysis,

as described elsewhere.19 The pH measurements were performed using a DKK-TOA

HM-25R pH meter.

Materials All solvents and reagents were of the highest qualities available and were used

as received without further purification. Na2[Ni II(dcpdt)2]•4H2O (dcpdt = 5,6-

dicyanopyrazine-2,3-dithiolate)10a and Na2[Ni II(qdt)2]•6H2O (qdt = quinoxline-2,3-

dithiolate)10b were synthesized as previously described. The compounds 3,7-

bis(dimethylamino)-phenothiazin-5-��-Hikotaro

Co., Ltd.), and potassium hexacyanoferrate(III) (K3[FeIII (CN)6��

Inc.) were used as received.

53

Determination of the turnover frequencies (TOF’s) The TOF in the electrochemical hydrogen evolution reaction (HER) can be

determined by using the following eq. 1,20

�� = �7[9:�;<=;>�"]"@ A B<CDE[��]F (1)

where G[HB�I� I��"]"@ is the diffusion coefficient of [NiII(dcpdt)2]2− (aproximately

obtained 2.8 × 10-6 [cm2 s-1��ic is the maximum current at the catalytic

peak, n is the number of electrons (2 electrons), F is the Faraday constant (9.6485

× 104 [C mol-1]), A is the surface area of the electrode (0.152 × 3.14 = 7.07 × 10-2

[cm2]), and [cat] is the concentration of [NiII(dcpdt)2]2− in solution (5.0 × 10-7 [mol

cm-3])). The kobs (i.e., TOFmax) were calculated as TOFmax = 11.7, 6.7 and 2.1 s-1 at

pH = 4, 5 and 6, respectively.

Estimation of the diffusion coefficients (D’s) Attempts to determine the diffusion coefficient (D) of [Ni(dcpdt)2]2− failed due

to the strong adsorption of the catalyst over the electrode surfaces while its

polarization under all conditions examined, regardless of the choice of solvent, such

as water, acetonitrile, DMF and so on. Because of such problematic situation, it was

decided to adopt the D value estimated for [NiII(qdt)2]2− as a D value approximated

for [Ni(dcpdt)2]2−. It is considered that this is a reasonable approach because of the

good resemblance of the two catalysts from the viewpoints of both the molecular size

and the dianionic nature of the systems, as depicted in the following structure

diagrams (Scheme 3).

Scheme 3. The structures of [NiII(dcpdt)2]2− (left) and [NiII(qdt)2]2− (right).

54

In the determination of the D value of [NiII(qdt)2]2−, an additional problem arose

due to its adsorption onto the glassy carbon electrode surfaces only upon the anodic

polarization performed for its aqueous solutions, even though the adsorption was

found to be negligible upon the polarization using its DMF solutions. Because of

these multiple problems, the D value of [NiII(dcpdt)2]2− was indirectly estimated by

the D(H2O) value of [NiII(qdt)2]2− using the observable D(DMF) value, as explained

below. In the approach in this thesis, the G"JKLMN values, i.e., O�� values defined in

eq. 2 (see below), were measured for the relevant systems including the reference

compounds (i.e., Methylene Blue and [FeIII (CN)6]3−). The measurements were

carried out using the BAS RRDE-3A Rotating Ring Disk Electrode systems and a

glassy carbon disk electrode (5 mmφ, 0.196 cm2). The O�� values measured are

defined by the so-call Levich equation, eq. 2,21

= 0.62P�Q�G"JKLMNRM"�= 0.62P�Q�O��R�

where is the limiting current, P is the number of electron (1 electron), Q is the

surface area of the electrode (0.252 × 3.14 = 0.196 [cm2]), � is the concentration of

the electroactive species (5.0 × 10-7 [mol cm-3]), K is the kinematic viscosity [cm2 s-

1], and R is the angular rotation rate of the electrode [rad s-1]. One can understand

that D is a solvent-independent value, while O�� is dependent on the choice of

solvent. Therefore, the O��ST� value can be estimated by adopting the observed O��GU�� value using the following equation (eq. 3).

O��I�ST� = O��GU�� × WK�"X K7YD� ZL� [�

The validity of this approach has also been confirmed by measuring the O�� values of [FeIII (CN)6]3− and Methylene Blue. Moreover, the D values for the Ni

compounds could also be benchmarked using the literature values of D’s reported for

these reference compounds, as summarized in Table 1.

(2)

(3)

55

[a] Determined in aqueous solutions. [b] Determined in DMF solutions. [c] Values not determined directly due to the adsorption of the compound over the electrode surfaces. [d] O��I�ST, [\�]^_�]L� =O��GU�, [\�]^_�]L� × WK�"X K7YD� ZL� [�

[e] Values not determined due to the low solubility of the compounds. [f] Calculated from the reported values.24 [g] Suggested to be approximately equal to [NiII(qdt)2]2−. [h] G��I�[\�]^_�]L� = Gàb�Uc� × {O��I�ST, [\�]^_�]L� O��ST,Uc�}f �� [i] G��I�[�g��\�[]fL� = Gàb�Uc� × {O��ST, [�g��\�[]fL� O��ST,Uc�}f �� [j] Used as the benchmark. G��I�Uc� = Gàb�Uc� × {O��ST,Uc� O��ST,Uc�}f ��

Compounds O��(H2O) [a] O��I(H2O) O��(DMF) [b] AK�"X K7YDh FL� [h

G��I

[cm2 s-1]

Gàb

[cm2 s-1] ref

[Ni II(dcpdt)2]2− − [c] − − [c] − 2.8×10-6 [g] − −

[Ni II(qdt)2]2− − [c] 2.2×10-4 [d] 2.2×10-4 0.99 [f] 2.8×10-6 [h] − −

[Fe(CN)6]3− 4.1×10-4 − − [e] − 7.2×10-6 [i] 7.6×10-6 22

Methylene Blue 3.8×10-4 − 3.6×10-4 1.03 6.3×10-6 [j] 6.3×10-6 [j] 23

Table 1. Diffusion coefficients (D’s) of [NiII(dcpdt)2]2−, [NiII(qdt)2]2−, [FeIII (CN)6]3− and

Methylene Blue.

56

As summarized in Table 1, the ratio of the kinematic viscosities of H2O and

DMF (i.e., WK�"X K7YD� ZL� [� = 0.99)24 was used to estimate the O��I�ST�

value of [NiII(qdt)2]2− (see the footnote d for Table 1). To the contrary, the

WK�"X K7YD� ZL� [� value could be experimentally determined as 1.03 using the

O��ST� and O��GU�� values observed for Methylene Blue in the

experiments, confirming the validity of the approach. All G��I’s were estimated

from the ratio between {O��ST�}f � using the Gàb value reported for

Methylene Blue23 as a benchmark (see also the footnote j for Table 1). Furthermore,

as summarized in Table 1, the validity of the method can be further confirmed by the

fact that the G��I value (7.2 × 10-6 cm2 s-1), estimated for [FeIII (CN)6]3− using the

above method, matches well with its Gàb value (7.6 × 10-6 cm2 s-1).22 On the basis

of these observations, the G��I(H2O) value of [NiII(qdt)2]2− (2.8 × 10-6 cm2 s-1) was

decided to be adopted as an approximated value of G��I for [NiII(dcpdt)2]2− in

order to estimate the TOF value for the HER.

57

Figure 1. (a,b) LSVs and the plots of the limiting current versus ω1/2.of 0.5 mM [NiII(qdt)2]2− in

DMF (a,b), 0.5 mM [FeIII (CN)6]3− in water (c,d), 0.5 mM Methylene Blue in water (e,f), and 0.5

mM Methylene Blue in DMF (g,h).

0 5 10 15 20-0.10

-0.08

-0.06

-0.04

-0.02

0.00

i / m

A

ω1/2

0 5 10 15 200.00

0.02

0.04

0.06

0.08

0.10

i / m

A

ω1/2

0 5 10 15 200.00

0.02

0.04

0.06

0.08

0.10

i / m

A

ω1/2

0 5 10 15 200.00

0.02

0.04

0.06

0.08

0.10

i / m

A

ω1/2

-1.0 -0.8 -0.6 -0.4

0

10

20

30

40

50

60

Potential / V vs. Fc/Fc +

Cur

rent

/ µA

-0.4 -0.2 0.0 0.2

0

10

20

30

40

50

60

70


Cur

rent

/ µA

-0.8 -0.6 -0.4 -0.2 0.0

-30

-20

-10

0


Cur

rent

/ µA

-0.4 -0.2 0.0 0.2 0.4

0

10

20

30

40

50

60

70


Cur

rent

/ µA

(a) (b)

(c) (d)

(e) (f)

(g) (h)

slope = -1.29 × 10-3

R2 = 0.99

xobs = 2.2 × 10-4

slope = 2.14 × 10-3

R2 = 0.99

xobs = 3.6 × 10-4

slope = 2.42 × 10-3

R2 = 0.99

xobs = 4.1 × 10-4

slope = 2.21 × 10-3

R2 = 1.00

xobs = 3.8 × 10-4

400 rpm900 rpm

1600 rpm2500 rpm3025 rpm

400 rpm900 rpm

1600 rpm2500 rpm

400 rpm900 rpm

1600 rpm2500 rpm3600 rpm

400 rpm900 rpm

1600 rpm2500 rpm3025 rpm

58

Determination of the Gibbs free energy of activation (∆G‡exp)

The Gibbs free energy of activation (∆G‡exp) of [Ni II(dcpdt)2]2− was determined

by using following eq. 4 (Eyring-Polanyi equation),25

� = ijklm exp��− qr‡tl � (4)

where � is the reaction rate constant (= kobs), u is the transmission coefficient

(≈ 1), �5 is the Boltzmann’s constant (= 1.380 × 10-23 [J K-1]), � is the absolute

temperature (298 [K]), ℎ is the Plank’s constant (6.626 × 10-34 [J s]) and � is the

gas constant (8.314 [J K-1 mol-1]).

The ∆G‡exp can be calculated as 16.0, 16.3 and 17.0 kcal mol-1 at pH = 4, 5 and

6, respectively.

59

Results and Discussion Figure 2 shows a Pourbaix diagram for [NiII(dcpdt)2]2–, showing the first oxidation

and the first reduction processes observed in the range pH = 3~6. The catalytic current

for HER was shown to be triggered by the first reduction process,10a even though the

subsequent electrode processes, which may be rapidly proceeded within the catalytic

cycles, are not observable, for they are likely to be buried under the catalytic current in

the same potential domain. As a result, it was taken the efforts to unveil possible electrode

processes that take place during the catalytic processes based on the DFT calculations, as

adopted by other researchers5d,6b,6c,8b,16d,17b,18 (see also Scheme 2 for details of the

estimation of redox potentials and pKa values). In the DFT calculations, the first oxidation

potential for the pH-independent [NiII(dcpdt)2]2‒/[Ni II(dcpdt)(dcpdt+•)]‒ couple,

observable at 0.39 V vs. SCE at pH = 5~6 (Figure 2), is used to benchmark all the

remaining redox potentials. Moreover, the value of pKa = 5.0 determined in the Pourbaix

diagram is adopted to benchmark the pKa values for the remainders.

Figure 2 A Pourbaix diagram showing the pH-dependent redox properties of [NiII(dcpdt)2]2‒,

where the diagram is drawn using the values reported in ref. 10a.

-1.5

-1

-0.5

0

0.5

1

3 4 5 6

Pot

entia

l / V

vs.

SC

E

pH

[NiII(dcpdt)2]2–

E = +0.39

ET (e-)

pKa = 5.0

Pot

entia

l (V

vs.

SC

E)

pH

[NiII(dcpdt)(dcpdtH)]–

[NiII(dcpdt)(dcpdt+•)]–

[NiII(dcpdt)(dcpdtH2)]–

(next PCET � [NiII(dcpdtH)(dcpdtH2)]–)

60

As briefly described,10a the most stable form computed for the first one-electron-

reduced product, i.e., [NiII(dcpdt)(dcpdtH2)]‒ (equivalent to 1e--2H+ species), possesses

two protons attached to one of the two pyrazine rings (see Figure 3). The DFT results

indicate that the favorable path yielding this species is the ET-PT pathway, in which

reduction of [NiII(dcpdt)(dcpdtH)]‒ (i.e., 0e--1H+ species) into [NiII(dcpdt)(dcpdtH)]2‒

(i.e., 1e--1H+ species) (Ecal = -0.71 V vs SCE) initially occurs and protonation of it into

[Ni II(dcpdt)(dcpdtH2)]‒ (i.e., 1e--2H+ species) (pKa,cal = 13) proceeds in order to complete

the overall PCET process. In other words, the PT-ET path can be ruled out because the

protonation of [NiII(dcpdt)(dcpdtH)]‒ (i.e., 0e--1H+ species) into [NiII(dcpdtH)2] (i.e., 0e-

-2H+ species) prior to its reduction is unfavorable due to a low pKa,cal value (2.7)

correlated with this process. Importantly, the DFT results clarify that the next PCET step

also proceeds at a close electrode potential. Similarly, an ET-PT step is a favorable path

affording the second PCET product [NiII(dcpdtH)(dcpdtH2)]‒ (i.e., 2e--3H+ species). As

shown in Figure 3, reduction of [NiII(dcpdt)(dcpdtH2)]‒ (i.e., 1e--2H+ species) into

Figure 3. Experimental and calculated thermodynamic parameters for the corresponding

reduction/protonation of [NiII(dcpdt)2]2‒. Each redox potential (E) is given in V vs. SCE. All

possible structures with all possible spin states have been optimized at B3P86/6-311+G(2d,p)

level of DFT with the solvation in water taken into consideration (C-PCM), and the lowest-energy

state has been adopted for each species. See Tables S2-S9 for detail.

+0e– +1e‒ +2e‒

+0H+

+1H+

+2H+

+3H+

1st PCET

Ecal = –0.71 V

+H+

+H+pKa,cal = 2.7

pKa,cal = 5.0pKa,exp = 5.0

+e–

Ecal = –0.10 V

2nd PCET

[Ni(dcpdt)2]2– (closed-shell singlet)

[Ni(dcpdt)(dcpdtH)]– (closed-shell singlet)

[Ni(dcpdtH)2] (closed-shell singlet)

[Ni(dcpdt)(dcpdtH)]2– (doublet)

+H+pKa,cal = 13

+e–

[Ni(dcpdt)(dcpdtH2)]– (doublet)

+H+pKa,cal = 3.3

[Ni(dcpdtH)(dcpdtH2)] (doublet)

Ecal = –0.79 V

+e–

Ecal = –0.55 V

+e–

[Ni(dcpdtH)(dcpdtH2)]– (closed-shell singlet)

+H+pKa,cal = 7.2

[Ni(dcpdt)(dcpdtH2)]2– (closed-shell singlet)

61

[Ni II(dcpdt)(dcpdtH2)]2‒ (i.e., 2e--2H+ species) (Ecal = -0.79 V) may occur prior to the

protonation of [NiII(dcpdt)(dcpdtH2)]2‒ (i.e., 2e--2H+ species) into

[Ni II(dcpdtH)(dcpdtH2)]‒ (i.e., 2e--3H+ species) (pKa,cal = 7.2). Similarly, the PT-ET path

is much less favorable because of the low pKa,cal value (3.3) correlated with the

protonation of the 1e--2H+ species into the 1e--3H+ species. The SOMO (singly occupied

molecular orbital), given in the first reduction product (1e--2H+), is localized over one of

the pyrazineditholate π* orbital (Table S6), as briefly reported elsewhere.10a On the other

hand, the second reduction product (2e--3H+) has a closed-shell singlet ground state with

the HOMO localized over the doubly protonated pyrazinedithiolate chelate (Table S9),

where the triplet for [NiII(dcpdtH)(dcpdtH2)]‒ is only 0.3 kcal/mol higher in energy than

the singlet. Interestingly, the second PCET product (2e--3H+) has a structure in which one

of the pyrazine rings has a bent geometry due to its sp3-hydridized amine nitrogen centre

(see Figure 4a), where the dihedral angle between the two C2N2 planes within the ring is

estimated as 158°. In other words, one of the pyrazines is considered as a 1,4-

dihydropyrazine moiety.26 For the later discussion, it must be noted that two pyrazine

rings have a planar geometry in either the non-reduced or one-electron-reduced system

(i.e., 0e--1H+ or 1e--2H+ species). The most remarkable finding in this study is that metal

hydride species, often considered as key intermediates required to promote catalysis of

HER (see above), can be given as higher-energy tautomers for the second PCET product

(2e--3H+). Figure 4 exemplifies six possible tautomers for the 2e--3H+ species realized in

the DFT calculations. The second lowest-energy tautomer (Figure 4b) possesses an

energy only 0.6 kcal/mol higher than that of the lowest-energy tautomer (Figure 4a). It

reveals that one of the protonation sites can be shifted from the pyrazine nitrogen donor

to one of the sulfur donors without any significant sacrifice in energy. Moreover, this

species can be transformed into several less thermodynamically favorable hydride species

by raising its energy by ca. 12-15 kcal/mol. This closed-shell singlet tautomer (Figure 4c)

has a square-planar geometry, favored for the d8-Ni(II) complexes with strong ligand

fields, where one of the four coordination sites is occupied by a hydride donor. In the

dangling pyrazinedithiolate, the unligated S atom remains deprotonated, while the two

pyrazine nitrogen donors are both singly protonated. The S donor, ligated in a position

trans to the hydride donor, has a bond length of Ni-S = 2.21 Å, which is only slightly

62

��

��

��

��

��

��

��

��

��

��

��

��

��

��

��

63

��

��

��‒�

��

��

��

��

��

tautomerization

+e–

+H+

+H++e–

-H2-H2

tautomerization

64

by a free energy change of 12.2 kcal/mol at pH = 3. However, at higher pH conditions,

the path clearly becomes thermodynamically unfavourable (e.g., ΔG = 20.4 kcal/mol at

pH = 9, so ΔG‡ > 20 kcal/mol is expected). However, it has been separately confirmed

that [NiII(dcpdt)2]2‒ behaves as an active catalyst for HER even at pH = 9 (see Figure 7).

Indeed, it is essential to promote the second ligand-centered PCET reduction step and its

tautomerization in order to form the hydride intermediate required to evolve H2.

Figure 6. Proposed mechanism for HER catalyzed by [NiII(dcpdt)2]2‒.

1st Ligand-centered PCET+e-,+H+

+e-,+H+

+0.6 kcal/mol

H2

∆G‡hydride = +16.6 kcal/mol

-8.0 kcal/mol

-1.1 kcal/mol

TShydride

-3.2 kcal/mol

+16.0 kcal/mol

Metal-centered PCET

+e-,+H+

2nd Ligand-centered PCET

H2, H+H+

-11.2 kcal/mol

∆G = +17.7 kcal/mol at pH = 7.0

*

65

Figure 7. Electrochemical H2 evolution catalyzed by Na2[Ni II(dcpdt)2]•4H2O (1 µM) during the

controlled potential electrolysis at -1.3 V vs. SCE in an aqueous borate buffer solution ��

pH 9.0, 12 mL) containing NaCl (0.1 M) under Ar atmosphere. The working, counter, and

reference electrodes were a GC rod, a Pt plate, and a SCE, respectively. Applied overpotential =

0.52 V. The results clearly indicate that [NiII(dcpdt)2]2- is an active catalyst for HER even at pH

=9.0. The gradual raise in the H2 evolution rate during the electrolysis is attributable to the gradual

deposition of the catalyst itself, as previously reported.10a As discussed in the report,10a this

catalyst tends to be absorbed over the GC surfaces even by simple dipping of the catalyst solution

without electrolysis. It was also reported that the absorbed species preserve the NiS4 core based

on the energy dispersive X-ray fluorescence spectroscopy (see ref. 10a).

0 10 20 30

0

10

20

30

40

TO

N

Electrolysis time / min0 10 20 30

0

2

4

6

8

10

12

H2

evol

ved

/ µL

Electrolysis time / min

66

The ��Figure 6) for the transformation of the second lowest-energy

tautomer (2e--3H+��Figure 4b) into a nickel hydride intermediate has been computed

by supposing them in an open-shell singlet state. As shown in Figure 6, its Gibbs free

energy of activation (ΔG‡hydride) is estimated as 16.6 kcal/mol. Very interestingly, this

ΔG‡hydride value is almost consistent with that for ΔG‡

exp (16.3 kcal/mol), which is

determined from a relatively low turnover frequency (TOF) of 6.7 s-1 at pH = 5, where its

TOF is estimated for the electrochemical HER by this catalyst (see Experimental Section

for details). The five-coordinate TS (TShydride) has a trigonal bipyramidal geometry with

the two SOMOs separately localized over the Ni and one of the pyrazine rings (see Figure

6 and Table S22). The Mulliken spin density on the Ni ion (ρNi = 0.64) indicates that it

has a formal oxidation state of Ni(III) rather than Ni(II), providing the formulation of

TShydride as [NiIII (H)(dcpdtH2–•)(dcpdt)]– rather than [NiII(H)(dcpdtH2)(dcpdt)]–.

Moreover, the formation of TShydride can be satisfactorily confirmed by carrying out

the intrinsic reaction coordinate (IRC) calculation (see Figure 8). In this calculation, while

transferring the position of proton from the S donor to the Ni ion, a concomitant

enlargement in the bent angle in the pyrazine ring based on the intramolecular PCET step

takes place, where the dihedral angle between the two C2N2 planes increases from 167°

(2e--3H+��Figure 4b) to 173° (TShydride).

IRC calculations also reveal that the subsequent step affords the distorted square

pyramidal hydridonickel species (2e--3H+, �� Figure 4e and Table S13), when

calculated for its open-shell singlet state (Figure 4). This metastable hydridonickel(III)

species (ρNi = 0.72), which possesses an energy 1.1 kcal/mol lower than TShydride, can be

further transformed into the most stable hydride intermediate (2e--3H+��Figure 4c),

as discussed above. It should be noted here that it cannot be completely ruled out H2

formation paths preceded by abstracting a proton from surrounding proton donors in the

bulk, such as acetic acid, oxonium ion (H3O+), etc.

67

��

��

��

��

��

��

68

It has been also confirmed that the direct H-H coupling among the N-H and S-H

groups within a simple dcpdt ligand in 2e--3H+, without forming a nickel hydride (see

Scheme 5), is an extremely high barrier path (ΔG‡ligand ��Figure 9

for its IRC), which can be thus abandoned.

Scheme 5. H2 evolution path without hydride formation.

��

��

�

�

�

��

��

��

�

��

��

��

�

�

�

��

��

�

�

��

��

��

��

��

�

�

�

��

��

��

�

��

��

��

��

�

�

�

��

��

�

�

��

��

∆G‡ligand = +36.8 kcal/mol

36.2 kcal/mol

0.6 kcal/molH2

+2e-,+2H+

69

��

��

��

��

��

��

70

As a consequence, one of the most favorable paths leading to evolve H2 is proposed

in Figure 10 (see also Scheme 6 for details). In this diagram, the relative free energy

change corresponds to the driving force of electron transfer calculated with respect to the

first ligand-based reduction (i.e., [NiII(dcpdt)(dcpdtH)]‒ (0e--1H+) + e‒ →

[Ni II(dcpdt)(dcpdtH)]2‒ (1e--1H+��Ecal = -0.71 V vs SCE). It clearly shows that formation

of the second PCET product (2e--3H+) is more or less thermodynamically favored at the

electrode potential of -0.71 V vs. SCE, and becomes increasingly favorable at more

cathodic potentials. Moreover, both the hydride formation and the H2 elimination

processes exhibit pH-independent behaviors (pH = 3, 5 and 7 in Figure 10). Importantly,

the H2 elimination (i.e., 2e--3H+(hydride) → [Ni II(dcpdt)(dcpdtH)]‒ (0e--1H+) + H2) is

exothermic by more than 8.0 kcal/mol. It is also noteworthy that either protonation path

(i.e., [NiII(dcpdt)(dcpdtH)]2‒ (1e--1H+) + H+ → [Ni II(dcpdt)(dcpdtH2)]‒ (1e--2H+) or

[Ni II(dcpdt)(dcpdtH2)]2‒ (2e--2H+) + H+ → [Ni II(dcpdtH)(dcpdtH2)]‒ (2e--3H+))

increases its thermodynamic driving force as the pH is decreased, consistent with the

higher catalytic current for HER at lower pH conditions, as previously reported.10a

Figure 10. Free energy diagrams for the catalytic pathways of [Ni II(dcpdt)2]2‒ for HER at each

pH condition. Relative free energy for half reactions corresponding to electron transfer processes

are calculated with respect to the 0e--1H+/1e--1H+ couple (Ecal = -0.71 V vs. SCE).

-1.0

-0.8

-0.6

-0.4

-0.2

0.0

0.2

0.4

0.6

0.8

Rel

ativ

e fr

ee e

nerg

y (e

V)

-20

-15

-10

-5

0

5

10

15

Rel

ativ

e fr

ee e

nerg

y (k

cal/m

ol)

-1.0

-0.8

-0.6

-0.4

-0.2

0.0

0.2

0.4

0.6

0.8

Rel

ativ

e fr

ee e

nerg

y (e

V)

-20

-15

-10

-5

0

5

10

15

Rel

ativ

e fr

ee e

nerg

y (k

cal/m

ol)

-1.0

-0.8

-0.6

-0.4

-0.2

0.0

0.2

0.4

0.6

0.8

-20

-15

-10

-5

0

5

10

15

+e-

1e--1H+

+H+

+e- +H+

2e--3H+

(S-protonated)

0e--1H+

+H2

pH 7

2e--3H+

(Hydride)pH 3

pH 5

Reference 0e --1H+/1e--1H+

(-0.71 V vs. SCE)

Reaction coordinate

Rel

ativ

e fr

ee e

nerg

y (e

V)

Relative free energy (kcal/m

ol)

0e--1H+

1e--2H+

2e--2H+ 2e--3H+

2e--3H+

(TS)

71

Scheme 6. Free energy changes relevant to the catalytic processes by [NiII(dcpdt)2]2– (M2–) at

each pH condition, as reported elsewhere.18

2H+ + 2e– → H2 E02H

+/H2 = –0.241 –0.0592 × pH (V vs. SCE)

∆G2H+

/H2 = –2F(E02H

+/H2 – E)

MH – + e– → MH 2– E0MH

–/MH

2–,cal = –0.707 (V vs. SCE)

∆GMH+

/MH = –F(E0MH

–/MH

2–,cal – E)

MH 2– + H+ → MH 2– pKa,MH

2–/MH2

–,cal = 13.0

∆GMH /MH2+ = –RTln(10) × (pKa,MH

2–/MH2

–,cal – pH)

MH 2– + e– → MH 2

2– E0MH2

–/MH2

2–,cal = –0.786 (V vs. SCE)

∆GMH2+

/MH2 = –F(E0MH2

–/MH2

2–,cal – E)

MH 22– + H+ → MH 3

– pKa,MH22–

/MH3–,cal = 7.21

∆GMH2 /MH3+ = –RTln(10) × (pKa,MH2

2–/MH3

–,cal – pH)

MH 3– → MH 3

–(hydride)

∆GMH3–/MH3

–(hydride) = 12.3 kcal/mol

MH 3–(hydride) → MH – + H2

∆G MH3–(hydride)/MH

–+H2 = ∆G2H

+/H2 – ∆GMH

–/MH

2– – ∆GMH2–

/MH2– – ∆GMH2

–/MH2

2–

– ∆GMH22–

/MH3– –∆GMH3

–/MH3

–(hydride)

72

Conclusions In this chapter, the DFT calculations are conducted to clarify the mechanism of

the HER by [NiII(dcpdt)2]2–. The results of the calculated reduction potentials and the

pKa values reveal that (i) the first reduction occurs as ligand-based PCET, consistent

with the experiment results, (ii) the second ligand-based PCET reduction

consecutively proceeds over the pyrazine moiety at a close electrode potential, which

affords the doubly-reduced triply protonated species of [NiII(dcpdt)2]2–. The most

favorable species generated by this process is [NiII(dcpdtH)(dcpdtH2)]‒, but the

hydride species can be formed via the intramolecular PCET from

[Ni II(dcpdtH)(dcpdtH2)]‒ as the uphill step. It should be emphasized that both

electron and proton transfers from pyrazine to nickel forms hydridonickel(III). The

activation energy of this process is estimated to be 16.6 kcal/mol by the IRC

calculation, which reasonably matches with the observed activation energy. This is a

quite rare example showing intramolecular proton/electron transfer after the ligand-

based PCET path (Scheme 7), as described in General Introduction of this thesis.

The direct hydrogen elimination path within the ligand moiety after the ligand-

based PCET step is also successfully clarified by DFT calculations, showing much

larger activation barrier required relative to that for the hydride path. It indicates that

the hydride intermediate is required to form even after the ligand-based PCET.

Scheme 7. Classification of this study: the ligand-based PCET leading to HER by [NiII(dcpdt)2]2–

by path (iii) (i.e., Ligand-based PCET followed by the formation of hydridonickel via

intramolecular proton transfer).

L(ne–)

L

+ne-,+H+

M ML(ne–)

M

+ne–,+H+

Ligand-based PCET

(i) Metal-ligand-based PCET H

H

+H+

H2

transform

nPath niii))

+H+

H2

(ii) HER from ligands

∆Gligand = 36.2 kcal/mol

∆Ghydride = 16.6 kcal/mol

73

The present chapter reveals the importance of having appropriate

electron/proton acceptor sites in close proximity to a catalytically active metal center.

This study for the first time points out that such consecutive ligand-based PCET

reduction processes may be utilized to finely tune the overpotential required to drive

HER by transition metal molecular catalysts.

74

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76

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77

Chapter 3: Ligand-based PCET Reduction in a

Heteroleptic Ni(bpy)(dithiolene) Electrocatalyst

Leading to a Lower Overpotential for Hydrogen

Evolution

Introduction Hydrogen is considered as one of the simplest forms of stable energy and has

attracted attention as the cleanest, sustainable energy which does not emit CO2 upon

combustion.1–6 The extremely high fuel-to-electricity conversion efficiency (ca.

70%) in fuel cells7 encourages researchers to pursue a possibility to adopt hydrogen

energy in the society. An important approach has been to develop fast and robust

catalysts for hydrogen evolution reaction (HER) together with oxygen evolution

reaction (OER) in order to achieve practically useful solar-driven or

electrochemically-driven systems for hydrogen generation from water.1–6 At the same

time, it is still needed to further gain the knowledge and skills in controlling the

molecular catalysis of HER using a larger variety of transition metal based

systems.5,8–12

In the past two decades, researchers attempting to explore artificial

hydrogenase13-15 mimics ��

are shown in Scheme 1a) in order to examine the catalytic activity of various

homoleptic ML2-type complexes with M = Fe,16-18 Co,19-22 Ni,23-31 Mo,32 Rh,33 and

W.34,35 it was also reported on the electrocatalytic activity of several NiL2-type

complexes, showing their unique reaction paths permitting the formation of hydride

intermediates via two consecutive ligand-based proton-coupled electron transfer

(PCET) processes.28,30,31

78

Scheme 1. (a) Examples of 1,2-dithiolene ligands whose metal complexes were previously

reported as molecular catalysts for HER. (b,c) Homoleptic and heteroleptic nickel(II) complexes

examined as the catalysts for HER in this study. (d) The proposed pathway for the catalysis of

Ni II(bpy)(dcbdt) for HER, where E or ET is the electron transfer process, and C or PT is the

protonation step.

In spite of the great progresses made so far in the studies of such dithiolene

metal complexes, it is noticed that the studies on the heteroleptic MLL’-type

compounds having a single dithiolene ligand are extremely rare. Only one report can

be found,36 in which the catalytic activity of a nickel(II) diphosphine dithiolene

complex for HER in the presence of trifluoroacetic acid as a proton source was

reported.36 The fact motivated us to explore the activity of other heteroleptic systems.

In this context, it was decided to focus on a NiII(bpy)(dcbdt), (bpy = 2,2’-��

dcbdt = 4,5-dicyanobenzene-2,3-dithiolate) depicted in Scheme 1b, in the hope to

examine the effect of having a bpy moiety as a simple electron reservoir to promote

(b) (c)

L =

[Ni(dcbdt)2]2-Ni(bpy)(dcbdt)

(a)

ET-PT (PCET)

(d)

79

the catalysis of HER. A homoleptic NiL2 complex having the same dithiolene ligand

was also prepared to compare its activity for HER as a control (Scheme 1c). Here it

is demonstrated that the finding that the bpy-based reduction in [NiI(bpy)(dcbdt)]–

causes a tremendous increase in the basicity at the d9 NiI center, resulting in a EECC′

route to evolve H2 (Scheme 1d) with its catalytic current flowing at much more

positive potential relative to the formal bpy reduction potential, where E and C denote

electron transfer and chemical process (i.e., proton transfer in this case), respectively,

and the prime notifies that it is the rate-determining step (RDS).

80


Materials All solvents and reagents were of the highest qualities available and were used as

received without further purification. 4,5-Dicyanobenzene-1,2-dithiol,37

(nBu4N)2[Ni(dcbdt)2] (dcbdt = 4,5-Dicyanobenzene-1,2-dithiolate),37 and Ni(bpy)(qdt)38

were synthesized as previously described.

Synthesis of Ni(bpy)(dcbdt)•H2O A solution of 4,5-Dicyanobenzene-1,2-dithiol (0.192 g, 1.0 mmol) and triethylamine

(0.213 g, 2.1 mmol) in methanol (300 mL) was slowly added to a solution of

Ni IICl2(bpy)39 (0.286 g, 1.0 mmol) in water (100 mL). After stirring this solution for 1h

at room temperature, the resulting mixture was filtered. The product was washed with

acetone, and recrystallized by acetonitrile to give a purple powder (yield: 0.33 g, 0.78

mmol, 77.6 %). 1H NMR (DMSO-d6 / TMS, ppm): σ 8.48-8.62 (m, 4H), 8.31 (t, J = 9

��-TOF MS: m/z = 426.97 [M + Na]+ (Calcd for

C18H10N4NaNiS2: 426.96) Anal. calcd for C18H10N4NiS2•H2��

N, 13.24. Foun��

Experimental Methods 1H NMR spectrum was acquired on a JEOL JNM-ESA 600 spectrometer. ESI-TOF

MS spectrum was recorded on a JEOL JMS-T100LP mass spectrometer. CV, LSV and

bulk electrolysis were performed on a BAS ALS 602DKM electrochemical analyzer. For

these experiments, a glassy carbon working electrode (3 mmφ), a platinum wire counter

electrode and an Ag/Ag+ reference electrode (0.249 V vs SCE) were employed, where n-

Bu4NPF6 (0.1 M) was used as a supporting electrolyte and all potentials reported are given

relative to the Fc/Fc+ couple (Fc/Fc+ = 0.155 vs SCE).

Computational Methods Density functional theory (DFT) calculations were performed using Gaussian 16

packages40 to understand the structural and spin-state candidates. The structures were

fully optimized using the B3P86 density functuonal41 with the effect of solvation in DMF

81

taken into consideration using the conductor-like polarizable continuum model (C-PCM)

method.41 The 6-31+G(d,p) basis set was applied to all atoms. The use of B3P86/6-

31+G(d,p) level of DFT shows good consistency with theoretical and experimental results

(see Figure 1 and Table 1), as described elsewhere.31,43 The redox potentials were

calculated by employing the experimentally determined reduction potentials of

Ni II(bpy)(qdt) (Figure 2) as the benchmark (see Scheme 2 for detail). Moreover, the

pKaDMF value of acetic acid (pKa

DMF = 13.5)45 was adopted as the benchmark to calculate

the pKa values of the possible intermediates (see also Scheme 2 for detail).

82

Scheme 2. Isodesminc reaction methods43,44 for calculating redox potentials and pKa values based

on the experimentally determined values for the reduction potential of NiII(bpy)(qdt) (first

reduction: –1.62 V vs. Fc/Fc+�� –2.33 V vs. Fc/Fc+�� -2,3-

�� Figure 2) and pKa for acetic acid (pKaDMF = 13.5), respectively. F is Faraday constant,

R is the gas constant, and T is temperature (298.15 K).

(1) Determination of the redox potentials for Ox + e– → Red

Ox + e– → Red ∆G0 = –FE0

(A) For the case of NiII(bpy)(dcbdt)

Benchmark for the first reduction process: charge change from 0 to –1

Ni II(bpy)(qdt) + e– → [Ni I(bpy)(qdt)]–

∆G0ref1 = –FE0

ref1 E0ref1 = –1.62 V vs. Fc/Fc+

Ni II(bpy)(dcbdt) + [NiI(bpy)(qdt)]– → [Ni I(bpy)(dcbdt)]– + NiII(bpy)(qdt)

∆G0r = G([Ni I(bpy)(dcbdt)]–)solv + G(Ni II(bpy)(qdt))solv

– G(Ni II(bpy)(dcbdt))solv – G([Ni I(bpy)(qdt)]–)solv

∆G0r = –∆G0

ref1 + ∆G0 = FE0ref1 – FE0

E0 = –∆G0r /F + E0

ref1

Benchmark for the second reduction process: charge change from –1 to –2

[Ni I(bpy)(qdt)]– + e– → [Ni I(bpy–•)(qdt)]2–

∆G0ref2 = –FE0


[Ni I(bpy)(dcbdt)]– + [NiI(bpy–•)(qdt)]2– → [Ni I(bpy–•)(dcbdt)]2– + [NiI(bpy)(qdt)]–

∆G0r = G([Ni I(bpy–•)(dcbdt)]2–)solv + G([Ni I(bpy)(qdt)]–)solv

– G([Ni I(bpy)(dcbdt)]–)solv – G([Ni I(bpy–•)(qdt)]2–)solv

∆G0r = –∆G0


E0 = –∆G0r /F + E0

ref2

83

(B) For the case of [NiII(dcbdt)2]2–

Benchmark for the reduction process

[Ni I(bpy)(qdt)]– + e– → [Ni I(bpy–•)(qdt)]2–

∆G0ref2 = –FE0


[Ni II(dcbdt)2]2– + [NiI(bpy–•)(qdt)]2– → [Ni II(dcbdt)(dcbdt–•)]3– + [NiI(bpy)(qdt)]–

∆G0r = G([Ni II(dcbdt)(dcbdt–•)]3–)solv + G([Ni I(bpy)(qdt)]–)solv

– G([Ni II(dcbdt)2]2–)solv – G([Ni I(bpy–•)(qdt)]2–)solv

∆G0r = –∆G0


E0 = –∆G0r /F + E0

ref2

[Ni III (H)(dcbdt)2]2– + [NiI(bpy–•)(qdt)]2– → [Ni II(H)(dcbdt)2]3– + [NiI(bpy)(qdt)]–

∆G0r = G([Ni II(H)(dcbdt)2]3–)solv + G([Ni I(bpy)(qdt)]–)solv

– G([Ni III (H)(dcbdt)2]2–)solv – G([Ni I(bpy–•)(qdt)]2–)solv

∆G0r = –∆G0


E0 = –∆G0r /F + E0

ref2

(2) Determination of pKa value for A + H+ → AH+

A + H+ → AH+ ∆G0 = RTln(10)pKa

CH3COO– + H+ → CH3COOH

∆G0ref = RTln(10)pKa,ref pKa,ref = 13.5

A + CH3COOH → AH+ + CH3COO–

∆G0r = G(AH+)solv + G(CH3COO–)solv – G(A)solv – G(CH3COOH)solv

∆G0r = RTln(10)pKa – RTln(10)pKa,ref

pKa = ∆G0r/RTln(10) + pKa,ref

84

Figure 1. (Black dotted line) UV-vis absorption spectrum for DMF solution of 0.1 mM

Ni II(bpy)(dcbdt). (Colored solid lines) Spectral features simulated based on the TD-DFT

calculations using several different functionals (B3P86, M06L, B3LYP, BP86 and ωB97XD),

where the structure of NiII(bpy)(dcbdt) in its closed-shell singlet state was optimized at each

functional/6-31+G(d,p) level of DFT with solvation in DMF taken into consideration (C-PCM).

The results reveal that B3P86 and B3LYP are suitable compared with other functionals.

300 400 500 600 700 8000.0

0.5

1.0

1.5

2.0

2.5

3.0 Norm

alized oscillator strengthA

bsor

banc

e

Wavelength / nm

ObservedB3P86M06L

B3LYPBP86

ωB97XD

85

Figure 2. A cyclic voltammogram (CV) for DMF solution of 0.5 mM NiII(bpy)(qdt) containing

0.1 M n-Bu4NPF6 at room temperature under Ar atmosphere, recorded at a sweep rate of 100

mVs-1. NiII(bpy)(qdt) shows two reversible reductions at –1.62 V (∆Ep = 64 mV) and –2.33 V vs.

Fc/Fc+ (∆Ep = 81 mV), respectively.

Table 1. A comparison of first and second reduction potentials of NiII(bpy)(dcbdt) between the

experimental and theoretical values. The latter values were calculated by DFT with several

different functionals (B3P86, M06L, B3LYP, BP86 and ωB97XD) using the same methods as

described in Scheme S1. The results reveal that most functionals are suitable except for BP86,

where B3P86 gives more reasonable results than B3LYP.

E1 / V vs.

Fc/Fc+ ∆ / V

E2 / V vs.

Fc/Fc+ ∆ / V

Observed -1.54 - -2.33 -

B3P86 -1.47 +0.07 -2.30 +0.03

M06L -1.49 +0.05 -2.32 +0.01

B3LYP -1.47 +0.07 -2.39 -0.06

BP86 -2.01 -0.47 -2.33 +0.00

ωB97XD -1.47 +0.07 -2.33 +0.00

-2.5-2.0-1.5-1.0-10

0

10

20

Cur

rent

/ µA


86

Results and Discussion

Electrocatalytic Properties Figures 3 and 4 shows the cyclic voltammograms (CVs) of NiII(bpy)(dcbdt) (Figure

3) and [NiII(dcbdt)]2– (Figure 4) recorded for their DMF solutions. In the absence of acid,

Ni II(bpy)(dcbdt) shows two reversible reductions at –1.54 V (∆Ep = 67 mV) and –2.33 V

vs. Fc/Fc+ (∆Ep = 75 mV), which are assigned as the Ni- and bpy-based reductions,

respectively (see Scheme 1d). These assignments are well supported by the DFT results

(see below). On the other hand, [NiII(dcbdt)2]2– exhibits a reversible reduction at –2.51 vs

Fc/Fc+ (∆Ep = 70 mV). This reduction is ascribed to a ligand-based one-electron reduction

to afford [NiII(dcbdt)(dcbdt‒•)]3–, where the radical charcter is equally delocalized over

the two dcbdt ligands. This is also supported by the DFT results (see below). For both Ni

catalysts, the catalytic current ascribable to HER is gradually raised as the concentration

of acetic acid is increased (see Figure 3 and 4).

Figure 3. CVs for DMF solutions of 0.5 mM NiII(bpy)(dcbdt) in the presence of 0-30 equivalents

of acetic acid (pKaDMF = 13.5, E0,solv

HA/H2 = –1.40 V vs. Fc/Fc+)45 containing 0.1 M tetra(n-

butyl)ammonium hexafluorophosphate (n-Bu4NPF6) at room temperature under Ar atmosphere,

recorded at a sweep rate of 100 mVs-1. The inset shows the icat/ip vs. (C0H+)1/2 plots under these

conditions.

0

20

40

60

80

Cur

rent

/ µA

-2.5-2.0-1.5-1.0

0

10


0.0 0.10

5

10

i cat

/ i p

(C0H+)1/2

30 eq20 eq15 eq10 eq5 eq0 eq

87

Figure 4. CVs for DMF solutions of 0.5 mM (n-Bu4N)2[Ni II(dcbdt)2] in the presence of 0-30

equivalents of acetic acid (pKaDMF = 13.5, E0,solv

HA/H2 = –1.40 V vs. Fc/Fc+)45 containing 0.1 M

tetra(n-butyl)ammonium hexafluorophosphate (n-Bu4NPF6) at room temperature under Ar

atmosphere, recorded at a sweep rate of 100 mVs-1. The inset shows the icat/ip vs. (C0H+)1/2 plots

under these conditions.

Obviously, the first reduction process by NiII(bpy)(dcbdt) is not likely to be coupled with

the HER by this catalyst. On the other hand, the HER by [NiII(dcbdt)2]2– is clearly

triggered by the first ligand-based reduction (see Figure 4).

The insets in Figures 3 and 4 show that the icat/ip value is linear to the square root of

the acid concentration ((C0H+)1/2) for both catalysts. The icat/ip is defined by the following

eq. 1:12,46,47

�� = 20.446�� 1� where icat denotes the catalytic peak current, estimated from the shoulder of each scan,

ip is the one-electron reduction peak current in the absense of acid, R is the gas constant,

0

20

40

60

80

Cur

rent

/ µA

0.0 0.10

5

10

i cat

/ i p

(C0H+)1/2

-2.5-2.0-1.5-1.0

0

10


30 eq25 eq20 eq15 eq10 eq5 eq0 eq

88

T is the absolute temperature, F is the Faraday constant, and ν is the scan rate. From

the linear dependence of icat/ip vs. (C0H+)1/2, kobs in eq. 1 can be defined by eq. 2:12,46,47 �� = �� 2�

where kglobal corresponds to the rate of HER by these catalysts.12,46,47 From the slope of

the icat/ip vs. (C0H+)1/2 plots, the kglobal values for these catalysts using acetic acid are

estimated, as summarized in Table 2.

The experiments were similarly carried out using three different proton sources,

triethylammonium chloride (Et3NHCl), chloroacetic acid, and benzoic acid. Figure

5a shows comparison of H2 evolution curvatures observed using these acids, together

with the results for acetic acid. The kglobal values, the Ecat/2 values, and the

overpotentials for HER (η's) for all the combinations are summarized in Table 2,

Table 2. The thermodynamic and kinetic parameters correlated with the HER catalyzed by

Ni II(bpy)(dcbdt) and [NiII(dcbdt)2]2- under each acidic condition.

Catalyst Ni II(bpy)(dcbdt) [Ni II(dcbdt)2]2-

Acid Et3NHCl Chloroacetic acid Benzoic acid Acetic acid Et3NHCl Acetic acid

pKa 9.2[a] 10.0[b] 12.2[a] 13.5[a] 9.2[a] 13.5[a]

EHA/H2 / V vs. Fc/Fc+ –1.15[a] –1.36[b] –1.32[a] –1.40[a] –1.15[a] –1.40[a]

Ecat/2 / V

vs. Fc/Fc+ –2.11[c] –2.15[c] –2.20[c] –2.24[c] –2.51[c] –2.50[c]

η / mV 960[d] 790[d] 890[d] 840[d] 1350[d] 1110[d]

kglobal / M-1s-1 (1.4 × 104)[e] 7.2 × 103 1.2 × 103 3.9 × 102 8.3 × 103 5.9 × 102

k1 / M-1s-1 (6.3 × 1010)[e] 6.4 × 109 3.5 × 107 5.5 × 105 8.3 × 103 5.9 × 102

k2 / M-1s-1 (1.4 × 104)[e] 7.2 × 103 1.2 × 103 3.9 × 102 – –

Figure 8 7 6 3 10 4

[a] Values reported in ref. 45 [b] Values reported in ref. 48 [c] Estimated from the CV conducted in the presence of 30 equivalents of each acid (15 mM). [d] Calculated from the equation "η = EHA/H2 – Ecat/2"12 [e] Estimated from the slope shown in Figure 5b.

89

Figure 5 (a) Acid dependence of linear sweep voltammograms (LSVs) for DMF solutions of 0.5

mM NiII(bpy)(dcbdt) in ther presence of 30 equivalents of Et3NHCl (pKaDMF = 9.2�� ),45

chloroaectic acid (pKaDMF ��48 benzoic acid (pKa

DMF �� 45 and acetic

acid (pKaDMF �� 45 containing 0.1 M n-Bu4NPF6 at room temperature under Ar

atmosphere, recorded at a sweep rate of 100 mVs-1. (b) Plots of log(k1) (blue) and log(k2) (green)

versus pKa of acids. The value of Et3NHCl (blank circle) was estimated from each slope, which

shows the linear free energy relationship (see also Figure 3 and 6-9).

-2.5-2.0-1.5-1.0

0

50

100

150

200C

urre

nt /

µA


8 9 10 11 12 13 140

5

10

15

log(

k)

pKa

(a)

(b)

Acetic acid

Benzoic acid

Chloroacetic acidEt3NHCl

Et3NHClChloroacetic acidBenzoic acidAcetic acidNo acid

90


of benzoic acid (pKaDMF = 12.2, E0,solv

HA/H2 = –1.32 V vs. Fc/Fc+)45 containing 0.1 M n-Bu4NPF6

at room temperature under Ar atmosphere, recorded at a sweep rate of 100 mVs-1. The inset shows

the icat/ip vs. (C0H+)1/2 plots under these conditions.


of chloroacetic acid (pKaDMF = 10.0, E0,solv

HA/H2 = –1.36 V vs. Fc/Fc+)48 containing 0.1 M n-

Bu4NPF6 at room temperature under Ar atmosphere, recorded at a sweep rate of 100 mVs-1. The

inset shows the icat/ip vs. (C0H+)1/2 plots under these conditions.

0.0 0.10

5

10

15

i cat

/ i p

(C0H+)1/2

-2.5-2.0-1.5-1.0

0

50

100C

urre

nt /

µA



-2.5-2.0-1.5-1.0

0

50

100

150

Cur

rent

/ µA


0.0 0.10

10

20

i cat

/ i p

(C0H+)1/2


91


of triethylammonium chloride (pKaDMF = 9.2, E0,solv

HA/H2 = –1.15 V vs. Fc/Fc+)45 containing 0.1

M n-Bu4NPF6 at room temperature under Ar atmosphere, recorded at a sweep rate of 100 mVs-1.

The inset shows the icat/ip vs. C0H+ plots under these conditions.


of triethylammonium chloride (pKaDMF = 9.2, E0,solv

HA/H2 = –1.15 V vs. Fc/Fc+)45 containing 0.1

M n-Bu4NPF6 at room temperature under Ar atmosphere, recorded at a sweep rate of 100 mVs-1.


-2.5-2.0-1.5-1.0

0

50

100

150

200C

urre

nt /

µA


0.00 0.01 0.020

10

20

30

i cat

/ i p

C0H+

-2.5-2.0-1.5-1.0

0

20

Cur

rent

/ µA


0 eq1 eq2 eq3 eq4 eq

92

Figure 10. CVs for DMF solutions of 0.5 mM (n-Bu4N)2[Ni II(dcbdt)2] in the presence of 0-30

equivalents of triethylammonium chloride (pKaDMF = 9.2, E0,solv

HA/H2 = –1.15 V vs. Fc/Fc+)45

containing 0.1 M n-Bu4NPF6 at room temperature under Ar atmosphere, recorded at a sweep rate

of 100 mVs-1.

0.0 0.10

10

20

30

i cat

/ i p

(C0H+)1/2

-2.5-2.0-1.5-1.0

0

50

100

150

Cur

rent

/ µA



93

Among various molecular catalysis routes defined by Savéant et al.,46 it is found

the EECC′ mechanism depicted in Scheme 3 is the most suitable one to explain the

electrode processes adopted for the HER by the present heteroleptic NiII(bpy)(dcbdt)

catalyst.

Scheme 3. The catalytic scheme of EECC mechanism,46 where O, P, Q and B correspond to

Ni II(bpy)(dcbdt), [NiI(bpy)(dcbdt)]‒, [Ni I(bpy‒•)(dcbdt)]2– and [NiII(H)(bpy)(dcbdt)]‒,

respectively.

The validity of this choice is also supported by the DFT results, which will be

discussed in a later section. Eqs. 3 and 4 correspond to the definition of kglobal and

Ecat/2 for this mechanism.46

w�� = x��1 0 x��

�� y1 0 x�x��z�3�

≈ x�

!�� = !)�+� 0 �� {||}1 0 x��

�� y1 0 x�x��z~�� 4�

≈ !)�+� 0 �� y1 0 x��x�z

The most remarkable fearture is that the onset potential as well as Ecat/2 for the HER

catalyzed by NiII(bpy)(dcbdt) shows an anodic shift of ca. 0.1~0.4 V with respect to

the second reduction potential (i.e., E0P/Q), as depicted in Figure 5a. A reasonable

O + e– P

P + e– Q

Q + H+ B

B + H+ O + H2

E0O/P

E0P/Q

k1

k2

94

interpretation is that the second term in eq. 4 has a positive value. This tendency

becomes more obvious as the pKa is decreased (see the onset potential, together with

Ecat/2 for the red line (i.e., Et3NHCl case) in Figure 5a, in which the catalytic current

starts to flow at ca. –2.0 V vs. Fc/Fc+). In order to observe such behaviors, a condition

of k2C0H+ << k1C0

H+ (i.e., k2 << k1) must be satisfied. The overall assessment is that

(i) the k1 step is a rapid process which follows the second ET process and these two

consecutive paths may be considered as a PCET process on the basis of the ET-PT

scheme, and (ii) the final k2 step is a relatively slow process consistent with the

definition of the EECC′ mechanism with the last step being the RDS.

Using eqs. 3 and 4, the rate constants k1 and k2 for the individual acids, except

for Et3NHCl (see below for its explanation), are estimated, as listed in Table 2.

Importantly, the logarithm of either k1 or k2 shows a linear dependence on pKa,

consistent with that both rates are limited by the proton abstraction from each acid

(see Figure 5b and Scheme 4). As discussed above, the present catalysis of HER falls

into a category for which the rate of HER is relatively slow and the diffusion of acid

into the electrode surface does not limit the catalytic cycle, as observed by the linear

dependence on icat/ip vs. (C0H+)1/2 (see Figure 3, 6 and 7). It is important to note here

that, in general, the pKa governs whether the rate of catalysis or diffusion of acid

limits the overall reaction rate of HER.47,49 It has been well documented that the

diffusion of acid tends to limit the overall reaction rate when the catalytic cycle is

rapid due to the low pKa condition. For such cases, the icat/ip value shows a linear

dependence on C0H+ (Figure 8). To the contrary, the catalytic cycle limits the overall

reaction rate when the catalytic process is slower relative to the diffusion of acid. In

the latter case, the icat/ip shows a linear dependence on (C0H+)1/2, which is exactly the

case adopted by NiII(bpy)(dcbdt) for most of the acids tested. However, the catalysis

using Et3NHCl having the lowest pKa adopts the former category. As supplied as

Figure 8, the icat/ip is linearly correlated with C0H+. Moreover, the non-catalytic wave

corresponding to the [NiI(bpy)(dcbdt)]–/[Ni I(bpy‒•)(dcbdt)]2– redox couple, is seen

after the flow of catalytic current (see Figure 9), indicating that the diffusion of acid

is not fast enough to flow catalytic current at ca. –2.4 V vs. Fc/Fc+.

95

Scheme 4. The H2 elimination path after formation of [NiI(bpy‒•)(dcbdt)]2–, which proceeds via

the consecutive protonation (PT) steps.

96

The mechanism of the HER by [NiII(dcbdt)2]2– was similarly analyzed on the

basis of the CV experiments, as depicted in Figure 11. In sharp contrast with

Ni II(bpy)(dcbdt), the onset potential and Ecat/2 for HER are little shifted upon

changing the pKa of acid with their potentials kept close to the first reduction potential

of [Ni II(dcbdt)2]2– observed in the absence of acid (E1/2 = –2.51 V vs. Fc/Fc+) (see

Table 2).These pKa-independent behaviors indicate that the k1 path is the RDS.46

Although an ambiguity remains in classification of the mechanism after the RDS, it

is assumed the EC′EC mechanism (Scheme 5) is a plausible one, as discussed below.

The DFT results are also in line with this selection (see below).

Figure 11 Acid dependence of LSVs for DMF solutions of 0.5 mM [Ni II(dcbdt)2]2- in the presence

of 30 equivalents of Et3NHCl (red),45 and acetic acid (green)45 under the conditions similar to

Figure 5a (see also Figure 4 and 10).

Scheme 5. The reaction scheme of ECEC mechanism,46 where P, Q, Q' and B correspond to

[Ni II(dcbdt)2]2–, [Ni II(dcbdt‒•)(dcbdt)]3–, [Ni III (-H)(dcbdt)2]2– and [NiII(-H)(dcbdt)2]3–,

respectively.

-2.5-2.0-1.5-1.0

0

50

100

150

200

Cur

rent

/ µA


No acidAcetic acidEt3NHCl

P + e– Q

Q + H+ Q'

Q' + e– B

B + H+ P + H2

E0P/Q

k1

E0Q'/B

k2

97

For this mechanism, under the conditions of k1C0H+ << k2C0

H+ (i.e., k1 << k2) and

E0P/Q < E0

Q'/B, the following eqs. 5 and 6 are satisfied.46

w�� = x�� 5� !�� = !) +�� 6�

To ensure the homogeneous nature of the HER by these catalysts, the so-called

"rinse test"25 was conducted. In these experiments, the glassy carbon electrode once

used for observing the electrocatalytic HER by each catalyst was taken out from the

catalysis solution, and was then sweeped under the same CV conditions using a fresh

electrolyte solution which does not contain any catalyst. As no catalytic current flows

at the rinse test for both catalysts (Figure 12-15), deposition of any active species,

such as heterogeneous materials, can be ruled out, although deposition of unidentified

less catalytically active species was not negligible when the fastest catalytic cycle

was promoted using Et3NHCl (Figure 16-17). It is also important to note that the H2

evolved during the electrocatalytic HER by either catalyst was qualitatively detected

by using the gas chromatographic technique (data not shown), even though the

quantitative factors remain unexplored due to inevitable deposition of the catalyst at

the prolonged electrolysis time.

98

Figure 12. A LSV (red) with use of a GC working electrode for DMF solution of 0.5 mM

Ni II(bpy)(dcbdt) in the presence of 0.1 M n-Bu4NPF6 and 15 mM acetic acid ��KaDMF =

13.5, E0,solvHA/H2 = –1.40 V vs. Fc/Fc+)45. The CV result of rinse test after the 1st scan (blue),

recorded after replacing the electrolysis solution with the same solution free of the catalyst. It

indicates that no materials are absorbed over the GC electrode after 1 scan of CV. All these

electrochemical measurements are recorded at a sweep rate of 100 mV/s, under Ar atmosphere at

room temperature.


Ni II(bpy)(dcbdt) in the presence of 0.1 M n-Bu4NPF6 and 15 mM benzoic acid ��KaDMF

= 12.2, E0,solvHA/H2 = –1.32 V vs. Fc/Fc+)45. The CV result of rinse test after the 1st scan (blue),




room temperature.

Blank0.5 mM Ni(bpy)(dcbdt) 1 sweepRinse test

-2.5-2.0-1.5-1.0

0

20

40

60

80

Cur

rent

/ µA



-2.5-2.0-1.5-1.0

0

50

100

Cur

rent

/ µA


99


Ni II(bpy)(dcbdt) in the presence of 0.1 M n-Bu4NPF6 and 15 mM chloroacetic acid ��

pKaDMF = 10.0, E0,solv

HA/H2 = –1.36 V vs. Fc/Fc+)48. The CV result of rinse test after the 1st scan

(blue), recorded after replacing the electrolysis solution with the same solution free of the catalyst.

It indicates that no materials are absorbed over the GC electrode after 1 scan of CV. All these


room temperature.

Figure 15. A LSV (red) with use of a GC working electrode for DMF solution of 0.5 mM (n-

Bu4N)2[Ni II(dcbdt)2] in the presence of 0.1 M n-Bu4NPF6 and 15 mM acetic acid ��KaDMF

= 13.5, E0,solvHA/H2 = –1.40 V vs. Fc/Fc+)45. The CV result of rinse test after the 1st scan (blue),




room temperature.

-2.5-2.0-1.5-1.0

0

20

40

60

80

Cur

rent

/ µA


Blank0.5 mM [Ni(dcbdt)2]2- 1 sweepRinse test

-2.5-2.0-1.5-1.0

0

50

100

150

Cur

rent

/ µA



100


Ni II(bpy)(dcbdt) in the presence of 0.1 M n-Bu4NPF6 and 15 mM triethylammonium chloride (15

��KaDMF = 9.2, E0,solv

HA/H2 = –1.15 V vs. Fc/Fc+)45. The CV result of rinse test after the 1st

scan (blue), recorded after replacing the electrolysis solution with the same solution free of the

catalyst. It reveals that the catalyst is partly absorbed over the GC electrode. All these


room temperature.

Figure 17. A LSV with use of a GC working electrode for DMF solution of [NiII(dcbdt)2]2- (0.5

mM) in the presence of n-Bu4NPF6 (0.1 M) and triethylammonium chloride ��KaDMF =

9.2, E0,solvHA/H2 = –1.15 V vs. Fc/Fc+)45. The CV result of rinse test after the 1st scan (blue),


reveals that the catalyst is partly absorbed over the GC electrode. All these electrochemical

measurements are recorded at a sweep rate of 100 mV/s, under Ar atmosphere at room

temperature.

-2.5-2.0-1.5-1.0

0

50

100

150

200

Cur

rent

/ µA



-2.5-2.0-1.5-1.0

0

50

100

150

200

Cur

rent

/ µA


Blank0.5 mM [Ni(dcbdt)2]2- 1 sweepRinse test

101

DFT Calculations Some important insights into the mechanism of the HER by NiII(bpy)(dcbdt)

and [NiII(dcbdt)2]2– were given by DFT calculations. As shown in Figure 18, the

initial reduction process for NiII(bpy)(dcbdt) is unambiguously assigned as a simple

one-electron reduction of the dissolved catalyst. The computed first reduction

potential (Ecal), which is assigned as the Ni(II/I) on the basis of the Mulliken spin

density at the Ni center (ρNi = 0.86) in the reduction product (see Figure 19), is almost

consistent with the observed potential (Eobs). A square scheme showing possible

reduction and protonation products can be also developed (Figure 18). The results

clearly indicate that the ET-PT path is the much more favorable for the subsequent

process, in which the bpy-based reduction proceeds as the second reduction to afford

[Ni I(bpy–•)(dcbdt)]2– (i.e., 2e–-0H+ �� Figure 20), Protonation of this product

further affords the two-electron-reduced singly protonated species

[Ni II(H)(bpy)(dcbdt)]– (i.e., 2e–-1H+ species). This hydridonickel(II) intermediate

has a square-planar geometry (see also Figure 22). The PT-ET path is largely

disfavored under the pKaDMF = 13.5 condition adopted in the above experiments

considering the pKa value of 2.3 for the protonation of [NiI(bpy)(dcbdt)]– (i.e., 1e–-

0H+ species). In the same manner, more protonation of the above hydridonickel(II)

species (i.e., 2e–-1H+ species) is largely disfavored due to the low pKa value (7.0)

computed for this path. Moreover, the calculations also reveal that the reduction of

this hydride intermediate can only be promoted at Ecal = –2.40 V, which is rather

cathodic relative to the potential where the catalytic current for the HER flows (ca. –

2.2 V, see Figure 5a).

102

��

��

��

��

��

103

��

��

��

��ρNi��

104

��

��

��

��ρNi��

��

105

��

��

106

��

��

107

��

��

��‒��

��

��

��

��

��ρNi��

��

��

��

��

�� ′��

��

��

��‒��

��

��

��

108

��

��

��

ρNi�≈��

��

��

109

��

��

��

�� ′� ��

�� ∆G = –14.76

��

�� ′��

��

��

110

Scheme 6. Free energy changes relevant to the catalytic processes by NiII(bpy)(dcbdt) (M ) under

each acidic condition, developed in the same manner as reported elsewhere 44b

2H+ + 2e– → H2 E02H

+/H2 = –1.40 (V vs. Fc/Fc+��

–1.32 (V vs. Fc/Fc+��

–1.15 (V vs. Fc/Fc+��3NH+)

∆G2H+

/H2 = –2F(E02H

+/H2 – E)

AH → A– + H+

pKaDMF = 13.5 (acetic acid)

12.2 (benzoic acid)

9.2 (Et3NH+)

∆GAH/A–H

+ = RTln(10) × pKaDMF

M + e– → M– E0M /M

–,cal = –1.470 (V vs. Fc/Fc+)

∆GM–/M

2– = –F(E0M /M

–,cal – E)

M– + e– → M2– E0M

–/M

2–,cal = –2.304 (V vs. Fc/Fc+)

∆GM–/M

2–= –F(E0M

–/M

2–,cal – E)

M2– + H+ → MH – pKa,M2–

/MH–,cal = 17.45

∆GM2–

/MH– = –RTln(10) × (pKa,M

2–/MH

–,cal – pKa

DMF)

MH – + H+ → M + H2

∆G MH–/M+H2 = ∆G2H

+/H2 – ∆GM

–/M

2– – ∆GM–/M

2– – ∆GM2–

/MH–

111

Figure 27. Free energy diagram for the catalytic pathway of NiII(bpy)(dcbdt) for HER. Relative

free energy for half reactions corresponding to electron transfer processes are calculated with

respect to the 1e--0H+/2e--0H+ couple (Ecal = -2.30 V vs. Fc/Fc+).

Figure 28. The acid dependence of free energy diagrams for the catalytic pathways of

Ni II(bpy)(dcbdt) for HER. Relative free energy for half reactions corresponding to electron

transfer processes are calculated with respect to the 1e--0H+/2e--0H+ couple (Ecal = -2.30 V vs.

Fc/Fc+).

-40

-30

-20

-10

0R

elat

ive

free

ene

rgy

(kca

l/mol

)

-2.0

-1.5

-1.0

-0.5

0.0

Relative free energy (eV

)

0.0

-19.23 -19.23

0e--0H+

1e--0H+ 2e--0H+

2e--1H+

0e--0H+ + H2

+e-

+H+

+e-

+H+

Reference 1e--0H+/2e--0H+

(-2.30 V vs. Fc/Fc+)

+H+

-3.88

1e--1H+

-15.772e--2H+

-24.653e--1H+

-41.70

+e-

+H+

-26.94

-50

-40

-30

-20

-10

0

Rel

ativ

e fr

ee e

nerg

y (k

cal/m

ol)

-2.5

-2.0

-1.5

-1.0

-0.5

0.0


)

-50

-40

-30

-20

-10

0

Rel

ativ

e fr

ee e

nerg

y (k

cal/m

ol)

-2.5

-2.0

-1.5

-1.0

-0.5

0.0


)

-50

-40

-30

-20

-10

0

Rel

ativ

e fr

ee e

nerg

y (k

cal/m

ol)

-2.5

-2.0

-1.5

-1.0

-0.5

0.0


)

0.0

-19.23 -19.23

AcOHPhCOOH

Et3NHCl

0e--0H+

1e--0H+ 2e--0H+

2e--1H+

0e--0H+ + H2

e-

H+e-

H+


(-2.30 V vs. Fc/Fc+)

-26.94-28.72-32.81

-41.70-45.39

-53.23

112

Figure 29. The acid dependence of free energy diagrams for the catalytic pathways of

[Ni II(dcbdt)2]2- for HER. Relative free energy for half reactions corresponding to electron transfer

processes are calculated with respect to the 0e--0H+/1e--0H+ couple (Ecal = -2.76 V vs. Fc/Fc+).

-80

-70

-60

-50

-40

-30

-20

-10

0

Rel

ativ

e fr

ee e

nerg

y (k

cal/m

ol)

-3.5

-3.0

-2.5

-2.0

-1.5

-1.0

-0.5

0.0


)

-80

-70

-60

-50

-40

-30

-20

-10

0

Rel

ativ

e fr

ee e

nerg

y (k

cal/m

ol)

-3.5

-3.0

-2.5

-2.0

-1.5

-1.0

-0.5

0.0


)

0e--0H+ 1e--0H+

1e--1H+

2e--1H+

0e--0H+ + H2

e-H+

e-

H+


(-2.76 V vs. Fc/Fc+)

AcOH

Et3NHCl

0.0 0.0

-8.30-14.17

-30.98-36.85

-44.30

-61.70

113

Based on the above results, possible schemes for the catalytic cycles of two

catalysts are proposed in Figure 30 As depicted in this figure, the pKa values of the

initially formed hydride intermediates can be estimated as pKa = 19-20, indicating

that the basicity of nickel centers is essentially similar to each other after the

formation of these metal hydrides. It is thus reasonable to consider that the essential

difference in the rate of proton abstraction is rather caused by the considerable

difference in the basicity of the precursor that interacts with the proton sources. In

the DFT results, it is noticed that the first reduction product for the homoleptic

catalyst, i.e., [NiII(dcbdt)(dcbdt‒•)]3–, has a SOMO spread over the two dcbdt ligands

with no hybridization of metal d orbitals (see Figure 24), indicating that electron

injection into this orbital has a relatively small contribution to raise the basicity of

the filled dz2 orbital at the Ni(II) d8 center. On the other hand, there are two SOMO’s

in the doubly reduced heteroleptic catalyst, i.e., [Ni I(bpy–•)(dcbdt)]2– with both

orbitals more or less mixed with the metal d orbitals (see Figure 20). The lower-

energy SOMO (–4.43 eV) corresponds to the metal-ligand antibonding couple with

relatively high contributions from two dithiolate sulfur 3p orbitals. Moreover, the

higher-energy SOMO (–3.36 eV) is derived from the π*(bpy) orbital with effective

contributions from the nitrogen 2p orbitals and small but non-negligible contributions

from the metal d-orbitals. It must be therefore emphasized in the end that the electron

filling in these two SOMO’s largely contribute to raise the basicity of the filled dz2

orbital due to their closer location with regard to the filled dz2 orbital. These well

rationalize the reason why the heteroleptic system can achieve substantial increase in

the proton abstraction rate required to promote the PCET pathways.

114

Figure 30. Proposed mechanisms of HER catalyzed by (a) Ni II(bpy)(dcbdt) and (b)

[Ni II(dcbdt)2]2–.

(a)

(b)

115

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116

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119

Concluding Remarks

Development of the energy conversion process from electricity to hydrogen or

other chemical has been widely studied to realize the clean and renewable energy

society. Inspired by the reaction centers of hydrogenases in nature, which show

similar catalytic activities for hydrogen evolution reaction (HER) to noble metal

heterogeneous catalysts, molecular catalysts for HER having 1st row metal ions have

been studied in the last decades. Among these studies, redox-active ligands including

dithiolenes were found to enhance the proton-coupled electron transfer (PCET)

process, which is essential to establish the effective molecular catalysts for HER.

However, in the previous studies of metal bis(dithiolene) complexes as the catalysts

for HER, the PCET triggered by the ligand-based reduction is rate-determining

during the catalysis, especially under neutral acidic conditions. In this context, the

author has attempted to develop nickel(II) dithiolene molecular catalysts for

electrochemical HER by promoting the ligand-based PCET reductions. Furthermore,

the mechanisms of HER were elucidated by electrochemical and computational

studies.

In Chapter 1, a nickel pyrazinedithiolate ([NiII(dcpdt)2]2‒�� -

dicyanopyrazine-2,3-dithiolate), having a dithiolene structure, is shown to serve as

an efficient molecular catalyst for HER in fully aqueous media. This catalyst shows

effectively low overpotentials for HER (330-400 mV at pH = 4-6). Moreover, the

turnover number of catalysis reaches 20000 over the 24-h electrolysis with a high

Faradaic efficiency of 92-100%. The electrochemical and DFT studies reveal that

diprotonated one-electron-reduced species forms at pH < 6.4 via ligand-based PCET

pathways, leading to electrocatalytic HER without applying highly negative potential

required to generate low-valent nickel intermediates.

Chapter 2 successfully clarifies the mechanism of HER catalyzed by

[Ni II(dcpdt)2]2– in water, which proceeds via formation of a square-planar

hydridonickel(II) intermediate given by unprecedented structural transformation of a

doubly-reduced triply-protonated species [NiII(dcpdtH2)(dcpdtH)]–, afforded as a

result of two consecutive ligand-based reductions of [Ni II(dcpdt)(dcpdtH)]– through

120

PCET pathways. DFT caluculations also indicate that the transition state (TS) during

HER, optimized at open-shell singlet state, forms via the intramolecular PCET from

pyrazine to nickel center. This is a rare example of catalyst exhibiting such a behavior.

In Chapter 3, a square-planar NiII(bpy)(dcbdt) hydrogen evolution catalyst is

shown to exhibit a substantial acceleration in the proton abstraction rate due to the

increased basicity at the filled Ni dz2 orbital after formation of [NiI(bpy–•)(dcbdt)]2–

via consecutive two one-electron reductions (bpy = 2,2’-�� -

dicyanobenzene-1,2-dithiolate). This catalyst adopts the EECC′ mechanism in which

the rate of the first protonation step is by far higher than that of the second step. The

DFT calculations reveal that the first and second reductions are correlated with the

electron injection into the metal-ligand anti-bonding and π*(bpy) orbitals,

respectively, where the latter orbital shows non-negligible hybridization with the

nickel d orbital. In addition, a homoleptic catalyst [Ni II(dcbdt)2]2– is shown to adopt

the EC′EC mechanism with the rate-determing step being a hydride forming step,

consistent with the largely delocalized nature of the injected electron over the two

dcbdt ligands (π*(dcbdt) orbital). This work demonstrates the importance of raising

the basicity of a metal d orbital relevant to proton abstraction in order to promote

PCET which significantly lowers the overpotential for H2 evolution.

The results demonstrated in this thesis are expected to provide important

strategies towards the molecular designs and development of highly effective

hydrogen evolution catalysts bearing redox-active ligands, which can be realized by

promoting ligand-based PCET pathways.

121

Acknowledgements

This thesis is the summary of the author’s studies obtained during April 2013 to

March 2019, at the Department of Chemistry, Faculty of Science, Kyushu University,

under the direction of Dr. Ken Sakai, Professor at Kyushu University.

The author expresses his sincere gratitude to Professor Ken Sakai and Dr. Kosei

Yamauchi for their significant guidance, continuous encouragement and valuable

discussions. The author had the opportunity to learn various important knowledge

and experimental skills, which have had a great impact on how he designed and

conducted the research topics summarized in this thesis. The author wishes to express

his sincere thanks to Dr. Hironobu Ozawa for his kind support for his research

activities in Sakai Laboratory and the Leading Graduate School Program.

The author is deeply grateful to Professor Andrew A. Gewirth and Professor

Thomas B. Rauchfuss at the University of Illinois at Urbana-Champaign for giving

him the valuable opportunities to study in their laboratories.

The author is also grateful to Professor Sharon Hammes-Schiffer and Dr. Mioy

Huynh at the University of Illinois at Urbana-Champaign for their kindness and

valuable discussion. The aouthor could learn a lot about computational chemistry

from them.

Acknowledgement is made to all the former and present members of Sakai

Laboratory for their valuable suggestions, heartfelt encouragements and friendship.

The author is much indebted for the financial support of Research Fellowship of

the Japan Society for the Promotion of the Science for Young Scientists (DC1). The

author is also much indebted for the financial support from the Advanced Graduate

Course on Molecular Systems for Devices at Kyushu University. The author is also

much indebted for the financial support from the International Institute for Carbon-

122

Neutral Energy Research (WPI-I2CNER), sponsored by the World Premier

International Research Center Initiative (WPI), MEXT, Japan.

Finally, the author wishes to offer his thanks to his family for their continuous,

warm-hearted encouragements and for their financial support.

Keita Koshiba

March, 2019

123

List of Publications

Chapter 1

“A Nickel Dithiolate Water Reduction Catalyst Providing Ligand-Based Proton-

Coupled Electron-Transfer Pathways”

Keita Koshiba, Kosei Yamauchi, and Ken Sakai

Angew. Chem. Int. Ed. 2017, 56, 4247-4251.

Chapter 2

“Consecutive Ligand-based PCET Processes Affording a Doubly Reduced Nickel

Pyrazinedithiolate which Transforms into a Metal Hydride Required to Evolve H2”


Dalton Trans., 2019, 48, 635-640.

Chapter 3

“Ligand-based PCET Reduction in an Heteroleptic Ni(bpy)(dcbdt) Electrocatalyst

Leading to a Lower Overpotential for Hydrogen Evolution”


Submitted to ChemElectroChem.

124

Other Publications

“電極材料、電極材料の製造方法および還元反応装置”

黄文しん、小柴慧太、古川晴一、宮島友博

特願 K2017-0119、出願日 2017年 11月 6日

“A family of molecular nickel hydrogen evolution catalysts providing tunable

overpotentials using ligand-centered proton-coupled electron transfer paths”

Yutaro Aimoto, Keita Koshiba, Kosei Yamauchi, and Ken Sakai

Chem. Commun. 2018, 54, 12820-12823.

Documents

Electrochemical and Theoretical Studies on Nickel