Electrocatalytic activity of cobalt tris(4-nitrophenyl)corrole for hydrogen evolution from water

Article abstract of DOI:10.1080/00958972.2019.1671588

Catalytic activities for hydrogen (H₂) evolution of cobalt 5,10,15-tris(phenyl)corrole (1) and 5,10,15-tris(4-nitrophenyl)corrole (2) (in homogeneous system was examined from acetic acid and water. In phosphate buffer solution (pH 7.0), the hyd

Full text of DOI:10.1080/00958972.2019.1671588

Journal of Coordination Chemistry
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Electrocatalytic activity of cobalt tris(4-
nitrophenyl)corrole for hydrogen evolution from
water
Hai Chen, Dong-Lan Huang, Md Sahadat Hossain, Guo-Tian Luo & Hai-Yang
Liu
To cite this article: Hai Chen, Dong-Lan Huang, Md Sahadat Hossain, Guo-Tian Luo & Hai-Yang
Liu (2019): Electrocatalytic activity of cobalt tris(4-nitrophenyl)corrole for hydrogen evolution from
water, Journal of Coordination Chemistry, DOI: 10.1080/00958972.2019.1671588
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JOURNAL OF COORDINATION CHEMISTRY
https://doi.org/10.1080/00958972.2019.1671588
Electrocatalytic activity of cobalt tris(4-nitrophenyl)corrole
for hydrogen evolution from water
Hai Chena,b, Dong-Lan Huangb, Md Sahadat Hossainb, Guo-Tian Luoa and
Hai-Yang Liub
aCollege of Chemistry and Chemical Engineering, Gannan Normal University, Ganzhou, China;
bDepartment of Chemistry, The Key Laboratory of Fuel Cell Technology of Guangdong Province,
South China University of Technology, Guangzhou, China
ABSTRACT
ARTICLE HISTORY
Received 29 January 2019
Accepted 5 September 2019
Catalytic activities for hydrogen (H2) evolution of cobalt 5,10,15-tri-
s(phenyl)corrole (1) and 5,10,15-tris(4-nitrophenyl)corrole (2) (in
homogeneous system was examined from acetic acid and water. In
phosphate buffer solution (pH 7.0), the hydrogen evolution turn-
over frequency (TOF) for 1 and 2 were 62 and 150hꢀ1 at the over-
potential of ꢀ1.53 V versus Ag/AgCl. In acetic acid media, hydrogen
evolution TOF for 1 and 2 were 56 and 127hꢀ1, respectively, at an
overpotential of 961 mV versus Ag/AgNO3. In the practical electroca-
talytic hydrogen evolution experiment at ꢀ1.45 V (vs. Ag/AgCl)
overpotential for 1 h, the TOF were 76 and 188hꢀ1 for 1 and 2,
respectively. No observable decomposition of the catalyst could be
observed after electrolysis for 72 h with a glassy carbon electrode.
KEYWORDS
Corrole; cobalt; hydrogen
evolution; electrocatalysis
1. Introduction
Development of renewable carbon-free energy is a major scientific and technological chal-
lenge to meet the energy needs for the future [1]. Hydrogen (H2) is considered as an ideal
energy carrier for our future power plants [2, 3]. One common approach to generate H2 is
CONTACT Guo-Tian Luo
luogt@gnnu.edu.cn; Hai-Yang Liu
chhyliu@scut.edu.cn
Supplemental data for this article can be accessed here.
ß 2019 Informa UK Limited, trading as Taylor & Francis Group

2
H. CHEN ET AL.
Scheme 1. Molecular structure of cobalt corroles.
the water splitting by electrochemical catalysis [4]. In nature, hydrogenase enzymes effi-
ciently catalyze hydrogen evolution. However, these enzymes are difficult to obtain and
their stabilities limit the application outside living systems [5]. Some platinum-based cata-
lysts can efficiently catalyze H2 production from water, but cannot be widely applied
because of its low natural abundance [6]. So, there is great demand in searching for hydro-
gen evolution catalysts of cheap metals. Many cheap metals such as iron [7], copper [8],
and cobalt [9] are capable of electrocatalytic hydrogen evolution.
Metallocorroles have been applied in catalytic oxidation [10], azacyclopropaneization
[11], and cycloaddition [12] reactions. However, the reports on the application of metal
corrole as electrocatalysts for H2 generation are so far limited to cobalt, copper, and man-
ganese complexes [13]. In 2014, Gross [14] reported that cobalt b-halogenated tris-
(pentafluorophenyl)corrole could efficiently electrocatalyze hydrogen evolution in
acetonitrile solutions. Cao [15] also used cobalt tris(pentafluorophenyl)corrole to catalyze
both water oxidation and proton reduction by coating it onto indium tin oxide (ITO) elec-
trodes in aqueous solutions, and found manganese corrole was unstable in the electroca-
talytic hydrogen evolution reaction (HER). Copper 5,15-pentafluorophenyl-10-(4-
nitrophenyl)corrole was also found active in electrocatalytic H2 evolution by using
trifluoroacetic acid as the proton source in acetonitrile [16]. Recently, we found that elec-
tron-deficient cobalt tris(ethoxycarbonyl)corrole exhibited good performance in the HER
[17]. Electron-withdrawing substituents can lower the overpotential significantly in the
cobalt corrole electrocatalyzed HER [15]. Similar observation was also found in the cata-
lytic oxygen reduction [18]. Nitrophenyl is an electron-withdrawing substituent that may
significantly tune chemical and physical properties of corrole derivatives [19]. Redox
potentials of cobalt corrole would be positively shifted (easier reduction and harder oxi-
dation) after attaching nitrophenyl substituents [20]. In 2001, Paolesse [21] reported the
first synthesis of tris(4-nitrophenyl)corrole. Since then, study on the electrochemical prop-
erty of its manganese complex [22], crystal structure of its copper complex [23], and its
(nitrosyl) iron complex [24] have also been reported. In 2017, Doctorovich [25] used cobalt
tris(4-nitrophenyl)corrole doped carbon nanotubes for photocatalytic and photo-electro-
catalytic hydrogen production. HER may be achieved at very low overpotential with high
efficiency, demonstrating cobalt tris(4-nitrophenyl)corrole is a promising catalyst for
hydrogen evolution. How about the electrocatalytic HER performance of this catalyst in
homogeneous solution? Here we report the electrocatalytic activity of cobalt 5,10,15-tri-
s(phenyl)corrole (1) and cobalt 5,10,15-tris(4-nitrophenyl)corrole (2) (Scheme 1) in HER.

JOURNAL OF COORDINATION CHEMISTRY
3
Water can be used as a proton source in acetonitrile-water (v/v ¼ 2/3) solvent system and
2 exhibits better performance than 1.
2. Experimental
2.1. Materials and measurements
All reagents were purchased from commercial sources and used without purification.
High resolution mass (HRMS) spectra were obtained with a Bruker Maxis impact mass
1
spectrometer incorporated with an electrospray ionization (ESI) source. H-NMR spectra
were recorded with a Bruker Avance III 400 MHz spectrometer in CDCl3 (d ¼ 7.26 ppm)
solvent. UV–visible spectra were measured with a Hitachi U-3010. Cyclic voltammo-
grams (CVs) were recorded with a CHI-660E electrochemical analyzer under nitrogen.
In the three-electrode cell in which a glassy carbon electrode (1 mm in diameter) was
used as the working electrode, saturated Ag/AgNO3 (in DMF solution) and Ag/AgCl (in
buffer solution) as the reference electrode, and a platinum wire as the auxiliary elec-
trode. The supporting electrolyte was 0.1 M [(n-Bu)4N]ClO4.
Controlled-potential electrolysis (CPE) experiments were carried out by employing
an airtight glass double-compartment cell which was separated by a glass frit in aque-
ous medium. The working compartment was fitted with a glassy carbon plate and a
Ag/AgCl reference electrode. The auxiliary compartment was fitted with a platinum
gauze electrode. The working compartment was filled with 40 mL of 0.25 M buffer
solution (acetonitrile/water (2:3 ¼ V/V) as solvent) containing 0.1 M KCl (supporting
electrolyte), the auxiliary compartment was also filled with 40 mL of phosphate buffer
solution (pH 7.0). Both compartments were sparged for 60 min with nitrogen. CPE
experiments were performed immediately after adding suitable concentration of com-
plex. After electrolysis, CH4 was added to the reactor as the internal standard, and the
gas mixture was analyzed by using an Agilent Technologies 7890A GC chromatograph.
2.2. Preparation of the cobalt complexes
Complexes 1 and 2: 5,10,15-tris(phenyl)corrole [26] (13.2 mg, 0.025 mmol), Co(OAc)2ꢁ4H2O
(30 mg, 0.12 mmol), and triphenylphosphine (PPh3) (50 mg, 0.191mmol) were dissolved in
ethanol (20 mL) and the solution was stirred for 1.5 h at room temperature.
Dichloromethane (DCM, 50 mL) was then added to the reaction mixture. The mixture was
washed three times with water, the organic phase was collected and the solvent was
removed by rotary evaporation to obtain the crude product. The product was purified by
column chromatography using DCM/Hex (V/V ¼ 2/1) as eluent. Recrystallization from
DCM/hexane afforded 1 as a red-brown solid (9 mg, yield 75%). Complex 2 was synthe-
sized by the same procedure using 5,10,15-tris(4-nitrophenyl)corrole [21]. HRMS spectra
of the complexes are given in the Supplementary materials (Figures S9 and S10).
1
Complex 1: UV-Vis (CH2Cl2) kmax (relative intensity) 384 (1.00), 560 (0.20). H-NMR
(400 MHz, CDCl3): 8.61 (d, 2H), 8.34 (d, 2H), 8.16–7.97 (m, 7H), 7.68–7.49 (m, 11H), 7.39
(d, 1H), 7.06 (t, 3H), 6.71 (t, 6H), 4.84–4.66 (m, 6H). HRMS, calcd for C55H38CoN4P:
845.2230, found for C55H38CoN4P: 845.2239.

4
H. CHEN ET AL.
1
Complex 2: UV-Vis (CH2Cl2) kmax (relative intensity) 372 (1.00), 575 (0.13). H-NMR
(400 MHz, CDCl3): 8.79 (d, 2H), 8.44 (m, 8H), 8.12 (d, 2H), 8.04 (d, 2H), 7.95 (d, 4H), 7.78
(d, 2H), 7.18 (t, 5H), 6.96 (t, 10H). HRMS, calcd for C55H35CoN7O6P: 980.1770, found for
C55H35CoN7O6P: 980.1791.
3. Results and discussion
3.1. Cyclic voltammetry
Cyclic voltammograms (CVs) of 1 and 2 in DMF are shown in Figure 1. The first (E1/2
¼
ꢀ0.51 V vs. Ag/AgNO3) and second (E1/2 ¼ ꢀ1.65 V vs. Ag/AgNO3) redox processes of
1 are reversible, assigned to the CoIII/II and CoII/I couples, respectively. Complex 2
exhibits three reversible redox peaks with E1/2 at ꢀ1.75, ꢀ1.02, and ꢀ1.53 V (vs. Ag/
II/I
ꢀ
AgNO3); these redox signals could be assigned to Co , nitrophenyl group redox [19a],
and CoIII/II couples, respectively.
The peak current corresponding to cobalt reduction potentials increases with the
scanning rate (Supplementary material, Figure S1). The peak current is linear to the
square root of the scanning speed. This indicates that the reduction is a diffusion con-
trolled process [2]. The CoII/I reduction current of 2 increased more quickly than 1
with the increase of scanning speed. This suggests diffusion rate of 2 is larger than
that of 1 in DMF.
3.2. Catalytic hydrogen evolution in DMF
To determine the electrocatalytic activity of 1 and 2, we recorded CVs in DMF in the
presence of acetic acid. For 1, adding different amounts of HOAc triggers catalytic cur-
rents at potentials near the CoII/I couple (Figure 2(a)). A systematic increase was
observed at ꢀ1.65 V versus Ag/AgNO3 with increasing acid concentrations from 0.00 to
2.20 mM as shown in Figure 2(a). This increase in current can be ascribed to catalytic
generation of H2 from acetic acid [27]. The onset of the catalytic wave transfers to
more anodic values with sequential increments of acid concentration. During addition
Figure 1. Cyclic voltammogram of 1 (a) and 2 (b) (1.15 mM) in DMF solution (0.1 M [n-Bu4N]ClO4,
scan rate of 100 mV/s).

JOURNAL OF COORDINATION CHEMISTRY
5
Figure 2. CVs of 1 (a) and 2 (b) (1.15 mM) in the absence of acid (black trace) and with different
concentrations of acid in DMF. Conditions: room temperature, 0.1 M [n-Bu4N]ClO4 as supporting
electrolyte, scan rate ¼ 100 mV/s GC working electrode (1 mm diameter), Pt counter electrode, and
Ag/AgNO3 reference electrode.
of HOAc a new hydrogen evolution peak can be observed near ꢀ1.91 V versus Ag/
AgNO3. Similar observations have been reported for other catalytic systems [28]. As in
-
these examples, hydrogen evolution may also form a cobalt-hydride H-CoIII intermedi-
ate, formed by the reduction of cobalt CoI specie. Thus there might be two hydrogen
evolution pathways A and B in the current catalytic system (Scheme 2). For 2 we can
see (Scheme 2) the mechanism of hydrogen production is similar to that of 1, but
since there is no new peak with increase of acidity, the hydrogen production path has
no B pathway, but we can see that at ꢀ1.02 V versus Ag/AgNO3 will change with the

6
H. CHEN ET AL.
Scheme 2. Possible mechanistic pathways for H2 evolution electrocatalyzed by complexes.
increase of acidity, in Figure 2(b), indicating that the p-nitrophenyl group is also
involved in hydrogen production. Combined with previous work [29], the electron
transfer in the catalytic process may be on the cobalt or on the ligand transfer.
Combined with the research of our group’s corrole complex electrocatalytic hydrogen
production, electron transfer is on cobalt because electron transfer is unstable on the
ligand [17].
A number of control experiments were also conducted under the same experimen-
tal conditions to confirm that 1 and 2 are responsible for the catalytic hydrogen evo-
lution, including the metal-free compound, Co(OAc)2ꢁ4H2O and the mixture of the free
ligand and Co(OAc)2ꢁ4H2O. As shown in Figures S3–S6 (Supplementary material), the
catalytic waves observed for the metal-free compound, Co(OAc)2ꢁ4H2O, or the mixtures
of the compound and Co(OAc)2ꢁ4H2O, did not match with 1 and 2.
Furthermore, bulk electrolyzes of a DMF solution with acetic acid at different
applied potential in a double-compartment cell were conducted to validate the elec-
trocatalytic activities of 1 and 2. Figure 3(a,c) shows the charge build up for bulk elec-
trolysis of solutions containing 1 and 2 in the presence of acetic acid. The charge
increased when the applied potential was increased. Another way to analyze good
catalytic efficiency is that catalysis could occur with relatively low overpotential. The
overpotential is defined by Equation (1) [30]; from Figure 3(a), we can see that when
the overpotential was 391 mV versus Ag/AgNO3, catalysis occurred. Further, when the
applied potential was ꢀ1.45 V versus Ag/AgNO3, the maximum charge reached 64.9
mC during 2 min of electrolysis,
Overpotential ¼ AppliedpotentialꢀE0
HA
(1)
E0 ¼ E0 ꢀð2:303RT=FÞpKaHA
HA
Hþ
where E0
is the standard potential for the solvated proton-dihydrogen couple and
Hþ
Ka,HA is the dissociation constant of acetic acid.

JOURNAL OF COORDINATION CHEMISTRY
7
Figure 3. Charge buildup vs. time from electrolysis of 1 (a) and 2 (c) (0.005 mM) in DMF (0.1 M [n-
Bu4N]ClO4) under different overpotentials. All data have been processed by blank subtraction.
Turnover frequency (mol H2/mol catalysts/h) for electrocatalytic hydrogen production by 1 (b) and
2 (d) (0.005 mM) under overpotentials (mV vs. Ag/AgNO3).
Presuming every catalyst molecule and every electron are used for the reduction of
protons, according to Equation (2) [31], we measured TOF values for the catalysts
reaching a maximum of 56 and 127 hꢀ1 at an overpotential of 961 mV (Supplementary
material, Equations (S1) and (S2)) for 1 (b) and 2 (d), respectively. Hence, the two com-
plexes have similar catalytic capacities, with 2 better than 1,
TOF ¼ DC=ðF ꢂ n1 ꢂ n2 ꢂ tÞ
(2)
where ᭝C is the charge from the catalyst solution during CPE minus the charge from
solution without catalyst during CPE, F is Faraday constant, n1 is the moles of elec-
trons required to generate 1 mol of H2, n2 is the moles of catalyst in solution, and t is
the duration of electrolysis.
3.3. Catalytic hydrogen evolution in aqueous media
We also studied catalytic hydrogen evolution in aqueous media, a more attractive
hydrogen power generation sustainable medium. CVs were carried out in 0.25 M phos-
phate buffers at different concentrations. At pH 7, the catalytic current was not appar-
ent until a potential of ꢀ1.31 V versus Ag/AgCl was attained in the absence of 2
(Figure 4(c)). With addition of 2, the onset of catalytic current was observed at about
ꢀ1.32 V versus Ag/AgCl, and the current increased significantly with increasing

8
H. CHEN ET AL.
Figure 4. CVs of 1 (a) and 2 (c) at different concentrations (pH ¼ 7.0). Conditions: 0.1 M KCl as
supporting electrolyte; scan rate, 100 mV sꢀ1; GC working electrode (1 mm in diameter); Pt counter
electrode; Ag/AgCl reference electrode. Plot of catalytic peak current vs. concentrations of 1 (b)
and 2 (d).
concentrations of 2 from 0.00 to 0.082 mM. The potential shifted to the positive value
of about 380 mV compared to that in the absence of 2, indicating that 2 can catalyze
the reduction of protons from water to H2. Catalytic hydrogen production can also be
achieved with 1 in neutral buffer (Figure 4(a,b)).
Figure 5(a,c) shows the bulk electrolysis in aqueous media at different applied
potential in a double-compartment cell. The total charge for bulk electrolysis of solu-
tions containing 1 and 2 increased when the applied potential was added. Based on
Equation (2), the TOF for 1 was 62 hꢀ1 at an overpotential of ꢀ1.53 V versus Ag/AgCl
(pH 7.0) (see Supplementary material, Equation (S3) and Figure 5(b)). Based on
Equation (2), the TOF for 2 was 150 hꢀ1 at an overpotential of ꢀ1.53 V versus Ag/AgCl
(pH 7.0) (see Supplementary material, Equation (S4) and Figure 5(d)). Comparing the
two results, 2 is a better catalyst than 1; also the TOF of 2 is higher than some
reported molecular catalysts for electrochemical hydrogen production from neutral
water. The TOF of a cobalt complex [32] in hydrogen production was only 5 hꢀ1 at
overpotential 390 mV. The turnover frequency for electrocatalytic water reduction was
1010 sꢀ1 at 1.0 V (vs. Ag/AgCl) of cobalt corrole [15] and some molecular catalysts [33]
such as Co-pentapyridine, molecular MoS2 catalyst, and Co-F8 were also lower than 2.
In summary, the performance of 2 for electrocatalytic hydrogen production is good in
the homogeneous phase.

JOURNAL OF COORDINATION CHEMISTRY
9
Figure 5. Charge buildup of 1 (a) and 2 (c) (0.005 lM) vs. different applied potentials in 0.25 M
phosphate buffer at pH 7.0. All data have deducted the blank. Turnover frequency (mol H2/mol cat-
alysts/h) for electrocatalytic hydrogen production by 1 (b) and 2 (d) (0.005 lM) under different
overpotentials (mV vs. SHE) in 0.25 M phosphate buffer (pH ¼ 7.0).
Figure 6. Complexes 1 (a) and 2 (b) GC traces after a 1 h controlled-potential electrolysis at
ꢀ1.45 V vs. Ag/AgCl of 0.005 lM complex 1 and 2 in 0.25 M phosphate buffer, pH 7.0. A standard
of CH4 was added for calibration purposes.
From the CPE experiments, the maximum charge was only 98 mC under ꢀ1.45 V
versus Ag/AgCl at pH 7.0 during 2 min of electrolysis in the absence of 1
(Supplementary material, Figure S6(a)). Upon addition of 1, charge reached to 183.6
mC due to the formation of H2. The evolved H2 was analyzed by gas chromatography.
Complex 1 (Figure 6(a)) gave 0.26 mL of H2 over an electrolysis period of 1 h with a

10
H. CHEN ET AL.
Faradaic efficiency of 90.7% for H2 and the TOF for 1 was 76 hꢀ1 under ꢀ1.45 V versus
Ag/AgCl. Complex 2 (6b) gave 0.64 mL of H2 over an electrolysis period of 1 h with a
Faradaic efficiency of 91.8% for H2 and the TOF for 2 was 188 hꢀ1 under ꢀ1.45 V ver-
sus Ag/AgCl.
3.4. Theoretical calculations
In order to explore the electronic structure of 1 and 2, density functional theory (DFT)
calculations were performed using the Gaussian 09 package [34]. The frontier molecu-
lar orbitals (FMOs) of 1 and 2 are shown in Scheme 3. The electron density of the
highest occupied molecular orbital (HOMO) and lowest unoccupied molecular orbital
(LUMO) were both distributed at the meso position as well as on the N atoms of pyr-
role rings, whereas they shifted towards the NO2 group when the para-hydrogen in
the meso-substituted phenyl group of corrole was substituted by NO2 group. The elec-
tron density of 2 has also decreased. This may be due to the stronger electronegativity
of the NO2 group, which increased the degree of electron migration to NO2 and
Scheme 3. FMOs of 1 and 2 and the ᭝E of LUMO and HOMO of complex (᭝E is the energy gap
of LUMO-HOMO).

JOURNAL OF COORDINATION CHEMISTRY
11
decreased the degree of electron overlap with other atoms. The push-pull effects
observed in the DFT studies are in agreement with the experimental data (red shift
and broadening of Soret band, alteration of redox potentials and HOMO-LUMO energy
gaps). In addition, the energy gap of LUMO and HOMO followed the order 1 > 2.
Since the ᭝E of 2 is smaller than that of 1, the ground state of 2 is easily excited. As
a result, 2 is more active than 1 in each chemical reaction [35]. Combined with our
experimental results, it can be seen that 2 is superior to 1 in the catalytic hydro-
gen production.
3.5. Catalytic stability of the complexes
To evaluate the catalytic durability of 1 and 2, we also operated extended CPE in
water while keeping the pH at 7.0 with 0.25 M buffer. A potential of ꢀ1.45 V versus
Ag/AgCl was employed for the measurement in order to ensure rapid turnover rate
during the electrolysis. It was shown (Supplementary material, Figure S8(c)), the total
of 415.1 C was passed for 2 during electrolysis. The catalyst affords a robust and essen-
tially linear charge build up over time, with no substantial loss during 72 h, and the
pH of the solution increased from 7.0 to 8.4, consistent with the accumulation of OHꢀ
by water reduction. Figure S7(a,b) (Supplementary material) displays the catalytic
charge and current for 1.
4. Conclusion
The two trisphenylcorrole cobalt complexes, 1 and 2, could be used as homogeneous
catalysts for electrochemical hydrogen evolution by using water as the proton source.
These complexes can electrocatalyze hydrogen evolution from acetic acid or aqueous
media (pH 7.0). Interestingly, in the practical electrocatalytic hydrogen evolution
experiment at ꢀ1.45 V versus Ag/AgCl overpotential for 1 h, 0.26 and 0.64 mL hydrogen
was produced for 1 and 2, respectively, showing the attachment of nitro groups on
the phenyl group impacts the performance of cobalt triphenylcorrole. The systematic
investigation of the effect of electronic structure on the electrocatalytic activity of
cobalt corrole in hydrogen evolution under the homogenous system continues in
our laboratory.
Disclosure statement
No potential conflict of interest was reported by the authors.
Funding
This research was funded by the National Natural Science Foundation of China (Nos. 21671068,
21161001) and Science and Technology Plan of Jiangxi Provincial Department of
Education (GJJ160941).

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H. CHEN ET AL.
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