A non-heme diiron complex for (Electro)catalytic reduction of dioxygen: Tuning the selectivity through electron delivery

Article abstract of DOI:10.1021/jacs.9b02011

In the oxygen reduction reaction (ORR) domain, the investigation of new homogeneous catalysts is a crucial step toward the full comprehension of the key structural and/or electronic factors that control catalytic efficiency and selectivity. Herein, we report a unique non-heme diiron complex that can act as a homogeneous ORR catalyst in acetonitrile solution. This iron(II) thiolate dinuclear complex, [Fe^II₂(LS)(LSH)] ([Fe₂^SH]⁺) (LS^2? = 2,2′-(2,2′-bipyridine-6,6′-diyl)bis(1,1-diphenylethanethiolate)) contains a thiol group in the metal coordination sphere. [Fe₂^SH]⁺ is an efficient ORR catalyst both in the presence of a one-electron reducing agent and under electrochemically assisted conditions. However, its selectivity is dependent on the electron delivery pathway; in particular, the process is selective for H₂O₂ production under chemical conditions (up to ～95%), whereas H₂O is the main product during electrocatalysis (less than ～10% H₂O₂). Based on computational work alongside the experimental data, a mechanistic proposal is discussed that rationalizes the selective and tunable reduction of dioxygen.

Full text of DOI:10.1021/jacs.9b02011

Subscriber access provided by UNIV OF LOUISIANA
Article
A non-heme diiron complex for (electro)catalytic reduction
of dioxygen: tuning the selectivity through electron-delivery
Lianke Wang, Marcello Gennari, Fabián G. Cantú Reinhard, Javier Gutiérrez, Adina Morozan,
Christian Philouze, Serhiy Demeshko, Vincent Artero, Franc Meyer, Sam P. de Visser, and Carole Duboc
J. Am. Chem. Soc., Just Accepted Manuscript • Publication Date (Web): 26 Apr 2019
Downloaded from http://pubs.acs.org on April 26, 2019
Just Accepted
“Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted
online prior to technical editing, formatting for publication and author proofing. The American Chemical
Society provides “Just Accepted” as a service to the research community to expedite the dissemination
of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in
full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully
peer reviewed, but should not be considered the official version of record. They are citable by the
Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore,
the “Just Accepted” Web site may not include all articles that will be published in the journal. After
a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web
site and published as an ASAP article. Note that technical editing may introduce minor changes
to the manuscript text and/or graphics which could affect content, and all legal disclaimers and
ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or
consequences arising from the use of information contained in these “Just Accepted” manuscripts.
is published by the American Chemical Society. 1155 Sixteenth Street N.W.,
Washington, DC 20036
Published by American Chemical Society. Copyright © American Chemical Society.
However, no copyright claim is made to original U.S. Government works, or works
produced by employees of any Commonwealth realm Crown government in the
course of their duties.

Page 1 of 24
Journal of the American Chemical Society
1
2
3
4
5
6
A non-heme diiron complex for (electro)catalytic reduction of dioxygen:
tuning the selectivity through electron-delivery
Lianke Wang,a Marcello Gennari,a,* Fabián G. Cantú Reinhard,b Javier Gutiérrez,a Adina
Morozan,c Christian Philouze,a Serhiy Demeshko,d Vincent Artero,c Franc Meyer,d Sam P. de
Visser,b,* Carole Duboca,*
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
a Univ. Grenoble Alpes, CNRS UMR 5250, DCM, F-38000 Grenoble, France
b Manchester Institute of Biotechnology and School of Chemical Engineering and Analytical
Science, The University of Manchester, 131 Princess Street, Manchester M1 7DN, United
Kingdom
c Université Grenoble Alpes, CNRS, CEA, Laboratoire de Chimie et Biologie des Métaux, F-
38000 Grenoble, France
d Universität Göttingen, Institut für Anorganische Chemie, Tammannstrasse 4, D-37077
Göttingen, Germany
* corresponding authors: carole.duboc@univ-grenoble-alpes.fr; marcello.gennari@univ-
grenoble-alpes.fr; sam.devisser@manchester.ac.uk
Abstract
In the Oxygen Reduction Reaction (ORR) domain, the investigation of new homogeneous
catalysts is a crucial step towards the full comprehension of the key structural and/or
electronic factors that control catalytic efficiency and selectivity. Herein, we report a unique
non-heme diiron complex that can act as a homogeneous ORR catalyst in acetonitrile
solution. This iron(II)-thiolate dinuclear complex, [FeII (LS)(LSH)] ([Fe2SH]+) (LS2- = 2,2’-
2
(2,2’-bipyridine-6,6’-diyl)bis(1,1-diphenylethanethiolate)) contains a thiol group in the metal
coordination sphere. [Fe2SH]+ is an efficient ORR catalyst both in the presence of a one-
electron reducing agent as well as under electrochemically assisted conditions. However, its
selectivity is dependent on the electron delivery pathway, in particular, the process is selective
for H2O2 production under chemical conditions (up to ~95%), whereas H2O is the main
product during electrocatalysis (less than ~10% H2O2). Based on computational work
alongside the experimental data, a mechanistic proposal is discussed that rationalizes the
selective and tunable reduction of dioxygen.
1
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 2 of 24
1
2
3
4
Introduction
5
6
7
8
The development of efficient, stable and noble metal-free catalysts for the O2 reduction
reaction (ORR) with selective production of either water or hydrogen peroxide remains one of
the central challenges in the energy conversion domain.1-2 Fuel cell design is focused on the
production of electricity based on the four-electron O2 reduction (to H2O) at the cathode,
associated with the oxidation of a fuel (such as H2 or MeOH) at the anode.3-5 On the other
hand, the two-electron O2 reduction is an interesting strategy to produce H2O2, which
represents a valuable commodity chemical and a versatile and clean oxidizing agent.6-8
Heterogeneous ORR catalysts have been mainly developed towards implementation into
exploitable devices for such applications.9-10 Among them, the most promising noble metal-
free materials are nitrogen-doped nanocarbon materials, containing molecularly-defined
active sites based on metal porphyrinic moieties.11-12 Regarding molecular homogeneous
catalysts, their study generally allows for a deeper comprehension of the ORR mechanism by
identifying the rate limiting steps for the reaction and the key structural and electronic factors
that control their activity and selectivity (two proton/two electron vs four proton/four electron
O2 reduction).1, 13 Such investigation should eventually lead to enhanced efficiency through
rational ligand design (e.g. by introduction of proton relays or Lewis acidic sites) for guiding
the bifurcation processes into the wanted reaction channel.
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
Regarding homogeneous ORR catalysts, the most common and well-studied ones are based
on iron or cobalt complexes embedded in macrocyclic heme-like ligands (such as porphyrins,
phthalocyanines, or corroles).14-21 Detailed electrochemical studies of a series of such
complexes evidenced how Fe-based catalysts are generally more selective for 4-electron O2
reduction than their Co analogues.1, 13, 22 This difference in reactivity is proposed to be a
consequence of the generation of high-valent metal-oxo species that are easier to generate for
hemes than for porphyrinic complexes based on other metal ions. Catalysts with non-heme
ligands have been much less explored, although some studies on copper-based complexes23-30
and, to a lesser extent, manganese-based systems have been reported.31-34 Even though in
several cases the nature of the acid and/or the presence of a proton relay were investigated on
such systems, the selectivity towards 4H+/4e- vs 2H+/2e- O2 reduction catalysis is still difficult
to predict.
For homogeneous catalysts, generally two ways to provide electrons have been adopted,
namely through the use of an appropriate reducing agent or an electrode (with application of a
suitable electrochemical potential). Both approaches have been widely tested but rarely with
the same catalyst. Using an Fe-based porphyrin system, Mayer et al. observed the same
2
ACS Paragon Plus Environment

Page 3 of 24
Journal of the American Chemical Society
1
2
3
4
5
6
7
8
9
efficiency and comparable selectivities for H2O2 production for both the chemically- and
electrochemically-driven processes (10(±3)% vs 10-20%).35 A mononuclear Cu complex
supported by a tetrapodal nitrogen-donor ligand was also reported and found to be selective
for 4H+/4e- O2 reduction catalysis in both cases. However, in that example electrocatalysis
was performed in protic solvent,36 while chemical reduction was carried out in acetone and
hence the environmental conditions affected the catalysis and selectivity.23
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
In the present work, we report the synthesis, characterization and reactivity of a dinuclear
non-heme iron complex, [FeII (LS)(LSH)] (LS2- = 2,2’-(2,2’-bipyridine-6,6’-diyl)bis(1,1-
2
diphenylethanethiolate),37 that acts as a homogeneous O2 reduction catalyst and displays a
change in selectivity under chemical versus electrochemical conditions. The process is
selective for H2O2 production when octamethylferrocene (Me8Fc) is used as reducing agent,
whereas H2O is dominant when catalysis is driven electrochemically, while the other main
experimental parameters (solvent, catalyst concentration, nature and concentration of the acid)
remained unchanged. With the aim of understanding how the selectivity could be fine-tuned
as a function of the different electron sources, a mechanistic investigation supported by
theoretical calculations has been performed. Surprisingly, even though the iron-oxygen
chemistry has been extensively investigated,38-44 our study reports the first non-heme iron-
based ORR catalyst. In addition, it provides insights on how to control the ORR selectivity in
separate pathways for 2-electron versus 4-electron reduction. Finally, comparison of the data
gained in this work with those previously reported for the isostructural Mn complex32,45
allowed us to provide insights into the role of the transition metal in ORR catalysis mediated
by such dimetallic complexes.
Results
Synthesis and solid-state properties of the dimercapto-bridged FeII dinuclear complex.
Under an inert atmosphere the LS2- ligand37 reacts with Fe(BF4)26H2O (1.2 equiv.) in THF to
yield a brown precipitate that was characterized as [FeII SH]+ or [FeII(LS)(LSH)]+ (Scheme 1).
2
Single crystal X-ray diffraction data collected on the isolated product reveal a dimercapto-
bridged dinuclear FeII complex [FeII (LS)(LSH)]BF4 ([Fe2SH]+), isostructural to the parent
2
MnII compound ([MnII (LS)(LSH)]BF4, [Mn2SH]+ reported previously.32
2
3
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 4 of 24
1
2
3
4
5
Feox
+
MeCN
Ph
Ph
6
7
N
FeIII
2
S
N
S
8
9
-
+
BF4
Ph
Ph
Ph
S
Ph
S
Ph
N
Ph
[FeMeCN +
]
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
H
S
Electrolysis
at -0.17 V
Fe(BF4)2

6H2O
N
N
N
FeII
FeII
Ph
S
+ 2 MeCN
- 2 MeCN
THF
N
MeCN
Ph
N
Ph
Ph
S-
-S
Ph
Ph
Ph
Ph
Ph
2+
Ph
Ph
S
Ph
LS2-
[Fe2
]
SH +
S
N
S
N
FeII
Ph
FeII
Ph
S
N
N
Ph
SS 2+
Ph
[Fe2
]
Scheme 1. Synthetic routes for the different iron-thiolate complexes described in this work.
Similarly to [Mn2SH]+, the structure of the [Fe2SH]+ cation (shown in Figure 1a, selected bond
distances and angles are reported in the Supporting Information) displays: (i) a planar {Fe2S2}
diamond core with deviations from the Fe1S2Fe2S3 plane of ~0.010 Å, (ii) two
distinguishable metal sites, each FeII center being pentacoordinated and surrounded by an
N2S3 donor set in a distorted trigonal bipyramidal environment (5 = 0.59 for Fe1 and 5 =
0.68 for Fe2),46 (iii) no direct metal-metal interaction, with a metal…metal distance of
3.1107(7) Å, and most importantly (iv) the unusual presence of one metal-bound thiol.47-50
The protonation of one thiolate is attested by the difference between the Fe1-S1 and Fe2-S4
bond lengths (2.3566(11) and 2.5387(11) Å, respectively) that gives evidence of the
protonation at S4. A similar ~0.2 Å discrepancy has been previously observed between the
two Mn-Sterminal distances in [Mn2SH]+ (Mn1−S1 = 2.4399(10) Å and Mn2−S51 = 2.6462(10)
Å). Thiolate protonation in [Fe2SH]+ is further confirmed by significant electron density
assigned to a hydrogen atom (H4) found on the difference Fourier map near S4 (S4 - H4 =
1.171(17) Å). The resulting S1…S4 distance (3.5940(10) Å) is compatible with the presence
of an intramolecular S4-H4…S1 hydrogen-bond, while in the parent Mn complex the S…S
distance is too long for such interaction (3.772(3) Å).
4
ACS Paragon Plus Environment

Page 5 of 24
Journal of the American Chemical Society
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
Figure 1. (a) Molecular structure of [Fe2SH]ClO41.5MeCN0.5Et2O determined by X-ray
crystallography with the thermal ellipsoids of the metal core drawn at 30% probability level.
All hydrogen atoms except H4, all anions and solvent molecules are omitted for clarity. (b)
Solid-state zero-field Mössbauer spectrum of [Fe2SH]+ recorded at 80 K.
The +2 oxidation state of both Fe ions is confirmed by Mössbauer spectroscopy (Figure 1b):
the zero-field spectrum of [Fe2SH]+ recorded on a solid sample at 80 K displays a single
doublet (δ = 0.87 mm.s-1 and ΔEQ = 3.76 mm.s-1) characteristic of two high-spin (S = 2) FeII
ions. Taken together with the charge balance of the complex (+1), these data are consistent
with the presence of one Fe-bound thiol in [Fe2SH]+. The two Fe centers are indistinguishable
by Mössbauer spectroscopy, reflecting the only small differences in their coordination
spheres, and suggesting that elongation of one Fe-S bond does not significantly alter the metal
ion’s electronic structure.
5
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 6 of 24
1
2
3
4
5
6
7
Solution and redox properties. In MeCN solution, [Fe2SH]+ remains dinuclear, as indicated
by (i) ESI-mass spectrometry, with a major peak at 1269.3 m/z (with isotope pattern
8
9
corresponding to [FeII (LS)(LSH)]+; Figure S1) and the absence of any peak originating from
2
1
mononuclear species, and (ii) the paramagnetic H NMR spectrum of [Fe2SH]+, displays no
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
large paramagnetic shift as observed in the spectrum of the parent mononuclear complex
[FeII(LS)Cl]-51 (Figure S2), consistent with an antiferromagnetically coupled dinuclear high
spin FeII complex.
Figure 2. Cyclic voltammograms of [Fe2SH]+ (~1.0 mM) before (black line) and after (blue
line) bulk electrolysis at -0.17 V vs Fc+/Fc in MeCN under argon atmosphere. Supporting
electrolyte: 0.1 M Bu4NClO4, GC working electrode ( = 3 mm), scan rate: 100 mV.s−1.
SH +
The cyclic voltammogram (CV, Figure 2) of [Fe2
]
displays an electrochemically
irreversible oxidation step (Epa = -0.19 V vs Fc+/Fc and Epc = -0.63 V), which is partially
reversible chemically (ipc/ipa ~0.6). On the basis of coulometric and spectroscopic studies
(vide infra), this redox system has been assigned to the two-electron metal-based oxidation of
the dinuclear complex (FeIIFeII→FeIIIFeIII) followed by chemical rearrangements. A mixture
of two different oxidized species is generated by bulk electrolysis carried out at -0.17 V vs
Fc+/Fc of a MeCN solution of [Fe2SH]+ (6.65 C of charge passed for 0.0368 mmol of initial
complex, corresponding to 0.94 electrons per Fe atom). By different spectroscopic techniques
1
(including UV-vis, H-NMR, and cw X-band EPR), they have been identified51 as: (i) a
mononuclear FeIII-thiolate complex, [FeIII(LS)(MeCN)]+ ([FeMeCN]+), and (ii) a dinuclear FeII-
6
ACS Paragon Plus Environment

Page 7 of 24
Journal of the American Chemical Society
1
2
3
4
disulfide complex, [FeII (LSSL)]2+ ([Fe2SS]2+, LSSL2- = disulfide form of the LS2- ligand,
2
5
6
7
Scheme 1). This assignment relies on our previous investigation showing that an identical
mixture is obtained when [Fe2SS]2+ is dissolved in MeCN. Note that in solid state, Feox can be
8
9
isolated in the dimeric form, [Fe2SS]2+.51 For simplicity, in the following paragraphs the
[FeMeCN]+ + [Fe2
]
mixture (~2:1 ratio)51 is designated Feox. Taken together, these results
SS 2+
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
imply that fast deprotonation accompanies the double electron transfer before the Feox
complex is formed.
A similar redox behavior is observed for the parent Mn-based dinuclear complex [Mn2SH]+,
II
which displays a ~180 mV anodic shift for the M2 /M2III redox system with respect to the Fe
SH +
complex (Epa = -0.19 V for Fe vs Epa = -0.01 V for Mn).32, 45 This shift indicates that [Fe2
]
is expected to be more easily oxidized by dioxygen with respect to the corresponding Mn
complex. For a previously described series of mononuclear MII-thiolate complexes, an even
larger cathodic shift (290 mV) was observed between the Fe and Mn derivatives.52
Homogeneous catalysis and electrocatalysis for dioxygen reduction. The capability of Feox
and [Fe2SH]+ to catalyze ORR was investigated in MeCN solution, in the presence of protons
and electrons. The 2,6-lutidinium tetrafluoroborate acid (LutHBF4) was employed as proton
source, whereas electrons were delivered by a chemical reducing agent (octamethylferrocene,
Me8Fc, E1/2 = -0.41 V vs Fc+/Fc, Figure S11) or alternatively by applying an adequate
SH +
electrochemical potential to a glassy carbon electrode. In the presence of LutH+, [Fe2
]
reacts with O2 to generate Feox (Figure S3). Note that this complex can also be prepared by
bulk electrolysis of [Fe2SH]+ (see above), or by dissolving the [Fe2SS](ClO4)2 solid in MeCN.51
Despite the fact that the redox potential of the Me8Fc+/Me8Fc couple is too high to reduce
Feox (Epc = -0.63 V, see above), Feox is efficiently reduced by Me8Fc in the presence of a
large excess (50 and 150 equiv.) of LutH+ to generate the oxygen reactive [Fe2SH]+ species
(see the UV-vis experiment in Figure S4). As expected, in the absence of protons, no reaction
is observed by mixing Feox and Me8Fc in MeCN solution under Ar atmosphere. The
protonation step that occurs after the reduction of Feox to generate [Fe2SH]+ is therefore
proposed to trigger the electron transfer process. Accordingly, in the CV of Feox the cathodic
wave is 60 mV shifted to a less negative potential (Epc = -0.57 V) after the addition of 150
equiv. of LutH+ (Figure S12).
As revealed by UV-vis absorption monitoring (Figure 3a, Table S3), the addition of a
catalytic amount of Feox (100 M) to an air-saturated (~1.6 mM.atm-1 O2) MeCN solution of
Me8Fc (2 mM, 20 equiv.) and LutH+ (15 mM, 150 equiv.) results in the progressive oxidation
7
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 8 of 24
1
2
3
4
5
6
7
8
9
of Me8Fc to Me8Fc+ (max = 750 nm,  = 420 M-1.cm-1), up to a ~75(±10)% amount. A similar
behaviour is observed when [Fe2SH]+ is employed as ORR catalyst (Figure S5), indicating that
both species are involved in the same catalytic cycle. In the absence of catalyst (and on the
timescale of catalysis) Me8Fc is not oxidized by O2. The absence of catalysis when the Fe
complexes [Fe(H2O)6](BF4)2 or [Fe(bpy)3](PF6)2 are employed, highlights the remarkable
beneficial effect of the coordination sphere provided by the thiolate-rich LS2- ligand. The non-
quantitative Me8Fc oxidation in the presence of Feox or [Fe2SH]+ can be rationalized in term of
partial catalyst deactivation during the ORR, most likely because of sulfur oxygenation in the
presence of in-situ generated H2O2 (Figure S9).
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
The amount of hydrogen peroxide produced in the reaction solution during the ORR catalysis
was quantified from a spectrophotometric assay using the Ti-TPyP reagent for the titration
(see Table S3).53-54 In the presence of Feox as catalyst, under the typical conditions of Figure
3, a 0.67(±0.03) mM concentration of H2O2 is detected after 10 min. If we take into account
the overall reaction stoichiometry (Eq. 1), this implies a 94(±4)% selectivity of the catalytic
process for H2O2 production.
O2 + 2 Me8Fc + 2 LutH+ → H2O2 + 2 Me8Fc+ + 2 Lut (Eq. 1)
Under these conditions, the ORR selectivity using Feox is thus comparable to that measured
with the parent Mn catalyst (82 (±2)%).
The kinetic profiles for the generation of Me8Fc+ catalyzed by Feox and the parent Mn catalyst
([Mn2SH]+) (Figure S6) indicate one order of magnitude faster apparent kinetics for the iron
system (TOFi Feox =8 (±1)103 h-1) compared to the manganese system (TOFi [Mn2SH]+ ≈ 4.0
(±0.4).102 h-1, TOFi = initial turnover frequency).
In the case of Feox, the TOFi is neither significantly sensitive to the relative concentration of
acid (from 5 to 15 mM) nor of reductant (from 2 to 4 mM), but the selectivity for H2O2 drops
when the concentration of reductant increases (46(±4)% with 4 mM of Me8Fc, Table S3).
8
ACS Paragon Plus Environment

Page 9 of 24
Journal of the American Chemical Society
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
Figure 3. (a) UV-vis spectral changes observed during ORR catalyzed by Feox in the presence
of LutH+ and Me8Fc in MeCN at 293 K: 2.0 mM Me8Fc, 15 mM LutHBF4, 0.1 mM (dimer)
Feox, air-saturated solution, 1 cm path length, t = 0.5 s. (b) Representative kinetic profiles
for the formation of Me8Fc+ at 750 nm in the presence of Feox ([Fe2SH]+ in Figure S5),
Fe(BF4)26H2O, or in the absence of catalyst (blank).
To test the ability of the air-stable Feox catalyst to serve as ORR electrocatalyst, CV
experiments of Feox (0.1 mM) were performed in the presence of a buffered LutH+/Lut MeCN
solution (15 mM, 150 equiv.) under saturated argon (Figure S14a) or oxygen atmosphere
(Figure 4a). Under O2 saturated atmosphere, an irreversible oxygen-reduction peak appears at
Ecat/2 = -0.50 V vs Fc+/Fc only in the presence of Feox. The ORR peak onset is only slightly
shifted (to less negative potentials) with respect to the cathodic peak of Feox under Ar
atmosphere (Epc = -0.60 V), but is ~5-fold more intense under the same experimental
conditions (see Figure 4a vs Figure S14a). These data unequivocally demonstrate the
electrocatalytic activity of Feox towards O2 reduction in MeCN solution. A rinse test confirms
that the activity arises from the iron catalyst in solution and is not due to any electrode-
adsorbed material (see Figure S13).
The standard potentials for 2H+/2e- and 4H+/4e- O2 reduction can be estimated in the buffered
LutH+/Lut MeCN solution (lutidinium, pKa = 13.92 in MeCN)55: E0(O2,4LutH+/2H2O) = 1.21
- 0.0592pKa ≈ 0.39 V vs Fc+/Fc, by using the method reported by Mayer et al.,1, 22 or
E0(O2,4LutH+/2H2O) = 2.038 - 0.0592pKa ≈ 1.21 V vs NHE, corresponding to 0.68 V vs
Fc+/Fc,56 by using the method reported by Nocera et al.57; E0(O2,2LutH+/H2O2) = 0.68 -
0.0592pKa ≈ -0.14 V vs Fc+/Fc (Mayer method).55 On the basis of the experimental mid-wave
potential of the catalytic wave (Ecat/2 = -0.50 V vs Fc+/Fc), the 4H+/4e- O2 reduction
overpotential (O2,4LutH+/2H2O) ( = Ecat/2-E0) is estimated in the range 0.89-1.18 V (while
the 2H+/2e- O2 reduction overpotential (O2,2LutH+/H2O2 is determined to 0.36 V). The
9
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 10 of 24
1
2
3
4
5
6
estimation of (O2,4LutH+/2H2O) is typical for homogeneous metal catalysts (mainly Fe- and
Co-based) for 4H+/4e- O2 reduction (0.52-0.96 V reported range).58
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
Figure 4. (a) Cyclic voltammograms (CV) on GC disk and (b) rotating ring-disk electrode
voltammograms (RRDEV) under oxygen-saturated atmosphere of a 15.0 mM LutH+ + 15.0
mM Lut solution in the absence (red lines) or presence (blue lines) of 0.1 mM (dimer) Feox.
MeCN 0.1 M Bu4NClO4, GC ( = 5 mm); scan rates: 100 mV.s−1 (CV), 5 mV.s−1 (RRDEV,
the response does not change within the range 2-10 mV.s–1); for RRDEV: rotation rate 1500
0.2V
rpm; ring current recorded with the Pt ring held at 0.2 V vs Fc+/Fc (for the blue curves, iring
has been corrected by subtracting iring0V, see Figure S16). The inset of (b) displays the
RRDEV-determined % of produced H2O2 and the number of electrons (n) consumed for
reduction of one molecule of O2 vs applied potential.
10
ACS Paragon Plus Environment

Page 11 of 24
Journal of the American Chemical Society
1
2
3
4
5
6
7
8
9
Rotating ring-disk electrode voltammetry (RRDEV) experiments were carried out for
estimating H2O2/H2O selectivity and quantifying the production of H2O2 under
electrochemical conditions (Figures 4b, S14-S15). The potential at the GC disk was scanned
through the catalytic region (from -0.4 V to -0.8 V vs Fc+/Fc), and the potential at the Pt ring
(Ering) was fixed at 0.2 (Figure 4b) or 0.43 V (Figure S15) vs Fc+/Fc57 in order to detect the
oxidation of the ORR-generated hydrogen peroxide.
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
Under argon atmosphere, the ring current observed in the presence of Feox and LutH+/Lut is
mainly attributed to the re-oxidation of [Fe2SH]+ formed on the disk, to regenerate Feox
(Figure S14). The ring current at 0.2 V vs Fc/Fc+ (blue trace in Figure 4b) in the presence of
Feox and LutH+/Lut under dioxygen atmosphere can be thus attributed to the oxidation of
H2O2 generated in the catalytic ORR process. Under these conditions, the Faradaic efficiency
for H2O2 production is ≤3.5% in the explored potential window, corresponding to ≳3.9
electrons (n) exchanged during the ORR process (see inset of Figure 4b). Similar results have
been found at Ering = 0.43 V vs Fc/Fc+: ≤12.5% H2O2, n ≳3.7 (Figure S15). These RRDE
studies, therefore, give evidence that the Feox catalyst is selective for water production, and
thus catalyzes almost exclusively 4H+/4e- O2 reduction under electrochemical conditions. A
bulk electrolysis carried out at -0.43 V vs Fc+/Fc on an O2-saturated acidic solution of Feox in
MeCN (0.1 mM complex, 15 mM LutH+), followed by Ti-TPyP reagent analysis, evidences
no formation of H2O2 in agreement with RRDE experiments.
Mechanistic investigations of the O2 reduction process
In support of the experimental observations, a series of density functional theory (DFT)
calculations were performed to gain insight into the mechanistic details for catalytic reduction
of dioxygen by [Fe2SH]+ (or Feox). We calculated individual structures along the proposed
reaction cycle of Scheme 2 (using two different models that give comparable results, see
Supporting Information, including Figures S17 – S22) and investigated reaction (free)
energies of dioxygen binding and activation.
The catalytic process is initiated by the reduction of Feox into [Fe2SH]+, the complex that does
interact with O2. The DFT-optimized structure of [Fe2SH]+ fully matches the X-ray structure,
and the high spin state of the FeII ions is well predicted. Specifically, the calculated distances
of Fe1-S1 and Fe2-S4 of 2.39 and 2.61 Å match the crystal structure bond lengths excellently
and indeed show elongation of the Fe-S bond for the thiol-ligated system. Note as well the
predicted hydrogen bonding interaction between the terminally bound thiolate and the thiol
(S-S = 3.582 Å) in the optimized structure of [Fe2SH]+ as observed experimentally. An almost
11
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 12 of 24
1
2
3
4
5
6
7
8
9
thermoneutral dioxygen binding step with G = 5.6 kcal mol–1 (E+ZPE = –2.2 kcal mol–1)
leads to the formation of the trans-1,2--peroxo dinuclear FeIII complex, [Fe2OO/SH]+ (via a
transient high-energy superoxo species, [Fe2OO°/SH]+, Figure S18 and Table S5). In contrast to
the corresponding DFT-calculated structure of the dinuclear MnIII-peroxo complex,30 the
diiron complex [Fe2OO/SH]+ is much less symmetric with an almost linear Fe–O–Fe angle of
173°, while the second oxygen atom is coordinated to one iron only (Figure 5). As such, the
structure of [Fe2OO/SH]+ resembles a side-on FeIII-peroxo linked to a second mononuclear FeIII
unit. Such -1:2-peroxo coordination mode has been already observed in heme-peroxo
copper models of the cytochrome c oxidase.59-61
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
Attempts to optimize either a more symmetrical structure for the trans-1,2--peroxo diiron
complex or a cis--1,2-peroxo coordination mode failed and converged back to the geometry
shown in Figure 5. The cis--1,2-peroxo mode is the exclusive coordination mode found in
non-heme FeIII peroxo-bridged complexes,62-64 in which one or more additional bridges (e.g.
carboxylate, phenolato) are often present to stabilize such species. Conversely, complexes
with unsupported peroxo-bridged dimers as in FeIII-tetraphenylporphyrins display a trans--
1,2 coordination mode as in the present case.65
We thus evaluated the capacity of the peroxo [Fe2OO/SH]+ complex to evolve either via an
internal proton transfer to form a hydroperoxo complex or via the breaking of the O-O bond
to generate a high-valent Fe-oxo species.
Concerning the intramolecular proton transfer pathway in [Fe2OO/SH]+ (Figure 5), a kinetic
barrier of 12.8 kcal mol–1 has been calculated to generate the diiron(III)-hydroperoxo complex
[Fe2OOH]+ (Figure S22). Regarding the cleavage of the O-O bond, even if two different
pathways were located (see SI for further details), the most favorable one corresponds to the
generation of dinuclear FeIV di-oxo intermediate [Fe2(µ-O)2SH]+ (Figure 5). However, the
energy barrier of 15.9 kcal mol–1 for this reaction is larger than the barrier predicted to
generate the protonated [Fe2OOH]+ peroxo complex. These calculations are thus consistent
with the fact that H2O2 can be selectively produced (in the presence of Me8Fc), through the
ox
reaction of [Fe2OOH]+ with an additional proton and the formation of Fe2 (Scheme 2,
pathway B).
12
ACS Paragon Plus Environment

Page 13 of 24
Journal of the American Chemical Society
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
Figure 5. Optimized geometries of [Fe2SH]+, [Fe2OO/SH]+, [Fe2(µ-O)2SH]+ and [Fe2OOH]+ in the
antiferromagnetically coupled singlet spin state with bond lengths in angstroms. For the sake
of simplicity, the peroxo intermediate is considered only in its protonated form. Reaction
energies (E+ZPE) and reaction free energies (G) at 298 K are given in kcal mol–1 for the
large model structures. Transition state energies (in kcal mol–1) are reported at
UB3LYP/BS1//UB3LYP/BS2 and contain zero-point and solvent corrections.
When electrons are delivered electrochemically, catalysis mainly occurs through a 4H+/4e- O2
reduction mechanism. In this case we propose that O2 activation involves the formation of
either [Fe2OO/SH]+ (or [Fe2OOH]+), which is then immediately reduced at the electrode before it
can react with an additional proton. Indeed in Scheme 2, pathway A requires the use of three
protons and two electrons to convert [Fe2OO/SH]+ into Feox with production of water. The full
prediction of this complicated mechanism involving four protons and four electrons in
13
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 14 of 24
1
2
3
4
5
6
7
8
9
sequential and/or coupled manner to produce H2O through an O-O bond cleavage is out of the
scope of the present study.
Pathways A and B are both exergonic, even if route B has a much lower driving force (GB =
–9.2 kcal mol–1 vs GA = –39.2 kcal mol–1). The fact that [Fe2SH]+ has a proton located in the
metal coordination sphere means that during the catalytic cycle there is competition between
dioxygen bond cleavage and proton transfer, whereby a possible pathway to H2O2 formation
becomes possible.
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
+
+
S
SH
Fe
S
Fe
N
SH
S
N
O
S
N
S
S
+ O2
N
N
O
Fe
Fe
N
N
N
SH +
OO/SH +
[Fe2
]
[Fe2
]
A
pathway
B
pathway
3 LutH+ + 2 MeCN + 2e
LutH+ + 2 MeCN
MeCN + Lut
LutH+ + 2e
2 H2O + 3 Lut
H2O2 + Lut
2+
+
S
S
- 2 MeCN
+ 2 MeCN
MeCN
Fe
N
S
N
Fe
Fe
2
N
N
S
N
S
N
S
SS 2+
[FeMeCN +
]
[Fe2
]
Feox
Scheme 2. Proposed 4H+/4e- (electrochemically driven) vs 2H+/2e- (chemical reducing agent)
O2 reduction catalysis by [Fe2SH]+ (or Feox). For the sake of simplicity, only the peroxo
[Fe2OO/SH]+ intermediate is depicted (and not the hydroperoxo [Fe2OOH]+ species).
Discussion and Conclusion
SH +
In this study, we report the reactivity of a unique dinuclear iron-thiolate complex, [Fe2
]
(and of its oxidized form, Feox), as an efficient ORR (electro)catalyst. Thiolate ligation has
been previously proposed to promote the activation of molecular oxygen,66 via (i)
stabilization of metal−peroxo intermediates,67-72 (ii) decrease of the activation barrier for O2
binding,52, 73 (iii) decrease of the metal redox potential,74 (iv) increase of the basicity of bound
substrates75 or (v) labilization of the trans sites.74 Our study confirms and further generalizes
14
ACS Paragon Plus Environment

Page 15 of 24
Journal of the American Chemical Society
1
2
3
4
5
6
7
8
9
this aptitude: the thiolate-rich N2S2-donor LS2- ligand, in the presence of FeII (this study) or
MnII (previously reported work)32, 45 promotes O2 activation at the coordinated transition
metal ion. Although the continuous regeneration of the initial complex [Fe2SH]+ during
catalytic turnover entails that sulfur-oxygenation is limited in the present system during the
ORR catalysis, this is not the case for other mononuclear Fe-thiolate complexes mainly
reported by the Limberg and Goldberg groups.76-80 Clearly, the thiolate groups of the LS2–
ligand system are much more difficult to oxidize than the cysteinate group used in the
Limberg and Goldberg systems. Indeed, our calculations indicate little or no spin density on
the thiolate groups.
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
A peculiar feature of these [M2SH]+ catalysts (M = Fe, Mn) is the presence of one thiol group
coordinated to the metal centers that can act as a proton relay during catalysis. This is a
unique example of a sulfur-based pendant acid in ORR catalysis, because most of them are O-
and/or N-based (carboxylic acids, ammonium, urea, etc.).1, 81-82 The nature of the proton relay
can still play a role as potential tunable agent of the selectivity; however no general trend can
be drawn up now. Interestingly, while Kojima et al observed an enhancement for 4H+/4e- O2
reduction catalysis in the presence of pendant pyridine close to a Cu site,28 Machan et al.
described a Schiff-based Mn catalyst with a pendant phenol that displays selective 2H+/2e- O2
reduction activity.33
The investigation of isostructural iron- and manganese-based catalysts evidences that even if
both display similar selectivity in the presence of a chemical reducing agent under similar
experimental conditions, [Fe2SH]+ is ~ 6 times more reactive than [Mn2SH]+, (on the basis of
the TOFi values determined in the presence of Me8Fc). This difference in efficiency can be
related to the reduction potential of the catalysts, [Fe2SH]+ being more easily oxidized than
[Mn2SH]+ (see above).
The most attractive properties of [Fe2SH]+ as an ORR catalyst are that: (i) this is the first
homogeneous ORR catalyst based on non-heme iron and (ii) its selectivity can be tuned as a
function on the mode of electron delivery (chemical vs electrochemical catalysis).
Although iron porphyrin systems have been largely investigated,1, 13-15, 83 this is not the case
for non-heme iron complexes. To our knowledge, only one family of Fe complexes supported
by non-heme ligands has been previously reported in the ORR domain.84 However, the
described tetrapyridyl mononuclear complexes only react with O2 stoichiometrically, and not
catalytically (the initial complex can be regenerated only in a two-step process). We believe
that the discovery of such non-heme iron catalyst for ORR catalysis will pave the way to new
structure-reactivity correlation studies encompassing both heme and non-heme compounds.
15
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 16 of 24
1
2
3
4
5
6
7
8
9
As hydrogen peroxide is generated in both cases (predominantly and up to ~95% or in less
than ~10% in chemical and electrochemical catalysis, respectively), it can be proposed that a
common intermediate, that is, the calculated iron-peroxo complex [Fe2OO/SH]+, is generated
during catalysis (Scheme 2). The reactivity of such a peroxo intermediate has been proposed
to be central for rationalizing the selectivity of homogeneous ORR catalysts. A study on a
series of CuII catalysts with tri- and tetrapodal nitrogen-donor ligands has evidenced that the
electronic and structural properties of the generated peroxo intermediates are critical for
tuning the selectivity of the ORR process.23
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
The tunable selectivity of the [Fe2SH]+ catalyst as a function of the electron delivering manner
implies that the peroxo [Fe2OO/SH]+ (or [Fe2OOH]+) intermediate can either (i) react with one
proton to produce H2O2 (with Me8Fc as reducing agent), or (ii) be reduced to generate H2O as
the main product after successive reduction and protonation steps (under electro-assisted
catalysis). These two pathways are supported by DFT calculations, which predict that both are
thermodynamically feasible. The production of H2O2 is further supported by the fact that the
formation of the hydroperoxo intermediate is thermodynamically and kinetically favored over
the generation of high valent FeIV di-oxo species arising from “direct” O-O bond rupture (see
Figure 5). Concerning H2O production, the O-O bond cleavage should thus occur after further
reduction of the (hydro)peroxo complex. The tuning of the selectivity can be interpreted in
terms of a competition between acid-base and redox processes with the [Fe2OO/SH]+ (or
[Fe2OOH]+) peroxo as the key intermediate. Indeed, when the concentration of the chemical
reductant is increased less H2O2 is generated (see Table S3), suggesting that the peroxo
intermediate needs to be in a more reducing environment to be reduced (and then to break the
O-O bond to generate H2O) before it gets protonated (to generate H2O2). When
electrochemically driven, the electron transfer kinetics for its reduction is likely faster because
it occurs in the reaction-diffusion layer at the immediate vicinity of the electrode, compared to
bulk solution.
The present study thus represents a substantial contribution towards the understanding of the
factors controlling catalytic efficiency and selectivity of homogeneous ORR catalysis, in
terms of ligand design, nature of transition metal(s), proton and electron delivery.
Supporting Information. Additional experimental data, crystallographic data, spectroscopic
characterizations, redox data and DFT calculations. This material is available free of charge
via the Internet at http://pubs.acs.org.
16
ACS Paragon Plus Environment

Page 17 of 24
Journal of the American Chemical Society
1
2
3
4
Acknowledgements
5
6
7
8
The authors gratefully acknowledge research support of this work by the Labex arcane (ANR-
11-LABX-003), the Agence National de la Recherche and the Deutsche
Forschungsgemeinschaft (grants no. ANR-16-CE92_0012_01 and DFG Me1313/14-1,
NiFeMim), the COST Action CM1305 (EcostBio) especially via the STSM 34962 and the
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
financial support of China Scholarship Council (L.W.). FGCR thanks the Conacyt Mexico for
a studentship.
References
1.
Homogeneous Molecular Catalysts and Electrocatalysts. Chem. Rev. 2018, 118, 2340-2391.
2. Gewirth, A. A.; Varnell, J. A.; DiAscro, A. M., Nonprecious Metal Catalysts for
Oxygen Reduction in Heterogeneous Aqueous Systems. Chem. Rev. 2018, 118, 2313-2339.
3. Adler, S. B., Factors Governing Oxygen Reduction in Solid Oxide Fuel Cell Cathodes.
Chem. Rev. 2004, 104, 4791-4843.
4. Gao, Z.; Mogni, L. V.; Miller, E. C.; Railsback, J. G.; Barnett, S. A., A perspective on
low-temperature solid oxide fuel cells. Energy Environ. Sci. 2016, 9, 1602-1644.
5. Shao, M.; Chang, Q.; Dodelet, J.-P.; Chenitz, R., Recent Advances in Electrocatalysts
for Oxygen Reduction Reaction. Chem. Rev. 2016, 116, 3594-3657.
6. Campos-Martin, J. M.; Blanco-Brieva, G.; Fierro, J. L. G., Hydrogen peroxide
Pegis, M. L.; Wise, C. F.; Martin, D. J.; Mayer, J. M., Oxygen Reduction by
synthesis: an outlook beyond the anthraquinone process. Angew. Chem., Int. Ed. 2006, 45,
6962-6984.
7.
cells in the presence of scandium ion. Energy Environ. Sci. 2015, 8, 1698-1701.
8. Yamada, Y.; Fukunishi, Y.; Yamazaki, S.-i.; Fukuzumi, S., Hydrogen peroxide as
Yamada, Y.; Yoneda, M.; Fukuzumi, S., High and robust performance of H2O2 fuel
sustainable fuel: electrocatalysts for production with a solar cell and decomposition with a
fuel cell. Chem. Commun. 2010, 46, 7334-7336.
9.
Chatterjee, S.; Sengupta, K.; Mondal, B.; Dey, S.; Dey, A., Factors Determining the
Rate and Selectivity of 4e-/4H+ Electrocatalytic Reduction of Dioxygen by Iron Porphyrin
Complexes. Acc. Chem. Res. 2017, 50, 1744-1753.
10.
Mukherjee, S.; Mukherjee, A.; Bhagi-Damodaran, A.; Mukherjee, M.; Lu, Y.; Dey,
A., A biosynthetic model of cytochrome c oxidase as an electrocatalyst for oxygen reduction.
Nature Commun. 2015, 6, 8467.
11.
Artero, V., Bioinspired catalytic materials for energy-relevant conversions. Nature
Energy 2017, 2, 17131.
12.
Jaouen, F.; Jones, D.; Coutard, N.; Artero, V.; Strasser, P.; Kucernak, A., Toward
Platinum Group Metal-Free Catalysts for Hydrogen/Air Proton-Exchange Membrane Fuel
Cells. Johnson Matthey Technology Review 2018, 62, 231-255.
13.
Fukuzumi, S.; Lee, Y.-M.; Nam, W., Mechanisms of Two-Electron versus Four-
Electron Reduction of Dioxygen Catalyzed by Earth-Abundant Metal Complexes.
ChemCatChem 2018, 10, 9-28.
17
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 18 of 24
1
2
3
4
5
6
14.
Zhang, W.; Lai, W.; Cao, R., Energy-Related Small Molecule Activation Reactions:
Oxygen Reduction and Hydrogen and Oxygen Evolution Reactions Catalyzed by Porphyrin-
and Corrole-Based Systems. Chem. Rev. 2017, 117, 3717-3797.
7
8
9
15.
Cytochrome c Oxidase, Myoglobin, and Hemoglobin. Chem. Rev. 2004, 104, 561-588.
16. Dogutan, D. K.; Stoian, S. A.; McGuire, R.; Schwalbe, M.; Teets, T. S.; Nocera, D.
Collman, J. P.; Boulatov, R.; Sunderland, C. J.; Fu, L., Functional Analogues of
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
G., Hangman corroles: efficient synthesis and oxygen reaction chemistry. J. Am. Chem. Soc.
2011, 133, 131-140.
17.
Activation. Acc. Chem. Res. 2007, 40, 543-553.
18. Yeh, C.-Y.; Chang, C. J.; Nocera, D. G., "Hangman" Porphyrins for the Assembly of a
Model Heme Water Channel. J. Am. Chem. Soc. 2001, 123, 1513-1514.
19. Fukuzumi, S.; Mochizuki, S.; Tanaka, T., Efficient reduction of dioxygen with
Rosenthal, J.; Nocera, D. G., Role of Proton-Coupled Electron Transfer in O-O Bond
ferrocene derivatives, catalyzed by metalloporphyrins in the presence of perchloric acid.
Inorg. Chem. 1989, 28, 2459-65.
20.
Fukuzumi, S.; Okamoto, K.; Gros, C. P.; Guilard, R., Mechanism of Four-Electron
Reduction of Dioxygen to Water by Ferrocene Derivatives in the Presence of Perchloric Acid
in Benzonitrile, Catalyzed by Cofacial Dicobalt Porphyrins. J. Am. Chem. Soc. 2004, 126,
10441-10449.
21.
Honda, T.; Kojima, T.; Fukuzumi, S., Proton-Coupled Electron-Transfer Reduction of
Dioxygen Catalyzed by a Saddle-Distorted Cobalt Phthalocyanine. J. Am. Chem. Soc. 2012,
134, 4196-4206.
22.
Wang, Y.-H.; Pegis, M. L.; Mayer, J. M.; Stahl, S. S., Molecular Cobalt Catalysts for
O2 Reduction: Low-Overpotential Production of H2O2 and Comparison with Iron-Based
Catalysts. J. Am. Chem. Soc. 2017, 139, 16458-16461.
23.
Kakuda, S.; Peterson, R. L.; Ohkubo, K.; Karlin, K. D.; Fukuzumi, S., Enhanced
Catalytic Four-Electron Dioxygen (O2) and Two-Electron Hydrogen Peroxide (H2O2)
Reduction with a Copper(II) Complex Possessing a Pendant Ligand Pivalamido Group. J. Am.
Chem. Soc. 2013, 135, 6513-6522.
24.
Tahsini, L.; Kotani, H.; Lee, Y.-M.; Cho, J.; Nam, W.; Karlin, K. D.; Fukuzumi, S.,
Electron-Transfer Reduction of Dinuclear Copper Peroxo and Bis-μ-oxo Complexes Leading
to the Catalytic Four-Electron Reduction of Dioxygen to Water. Chem. Eur. J. 2012, 18,
1084-1093.
25.
Fukuzumi, S.; Tahsini, L.; Lee, Y.-M.; Ohkubo, K.; Nam, W.; Karlin, K. D., Factors
That Control Catalytic Two- versus Four-Electron Reduction of Dioxygen by Copper
Complexes. J. Am. Chem. Soc. 2012, 134, 7025-7035.
26.
Fukuzumi, S.; Kotani, H.; Lucas, H. R.; Doi, K.; Suenobu, T.; Peterson, R. L.; Karlin,
K. D., Mononuclear Copper Complex-Catalyzed Four-Electron Reduction of Oxygen. J. Am.
Chem. Soc. 2010, 132, 6874-6875.
27.
transfer reactions of Cu-O2 and Cu2-O2 complexes. Coord. Chem. Rev. 2013, 257, 187-195.
28. Kotani, H.; Yagi, T.; Ishizuka, T.; Kojima, T., Enhancement of 4-electron O2
Fukuzumi, S.; Karlin, K. D., Kinetics and thermodynamics of formation and electron-
reduction by a Cu(II)-pyridylamine complex via protonation of a pendant pyridine in the
second coordination sphere in water. Chem. Commun. 2015, 51, 13385-13388.
18
ACS Paragon Plus Environment

Page 19 of 24
Journal of the American Chemical Society
1
2
3
4
5
6
7
8
9
29.
Adam, S. M.; Garcia-Bosch, I.; Schaefer, A. W.; Sharma, S. K.; Siegler, M. A.;
Solomon, E. I.; Karlin, K. D., Critical Aspects of Heme-Peroxo-Cu Complex Structure and
Nature of Proton Source Dictate Metal-Operoxo Breakage versus Reductive O-O Cleavage
Chemistry. J. Am. Chem. Soc. 2017, 139, 472-481.
30.
Schaefer, A. W.; Kieber-Emmons, M. T.; Adam, S. M.; Karlin, K. D.; Solomon, E. I.,
Phenol-induced O-O bond cleavage in a low-spin heme-peroxo-copper complex: Implications
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
for O2 reduction in heme-copper oxidases. J. Am. Chem. Soc. 2017, 139, 7958-7973.
31.
Evans, D. F.; Sheriff, T. S., The production of hydrogen peroxide from dioxygen and
hydroxylamine catalyzed by manganese complexes. J. Chem. Soc., Chem. Commun. 1985,
1407-8.
32.
Gennari, M.; Brazzolotto, D.; Pecaut, J.; Cherrier, M. V.; Pollock, C. J.; DeBeer, S.;
Retegan, M.; Pantazis, D. A.; Neese, F.; Rouzieres, M.; Clerac, R.; Duboc, C., Dioxygen
Activation and Catalytic Reduction to Hydrogen Peroxide by a Thiolate-Bridged
Dimanganese(II) Complex with a Pendant Thiol. J. Am. Chem. Soc. 2015, 137, 8644-8653.
33.
Hooe, S. L.; Rheingold, A. L.; Machan, C. W., Electrocatalytic Reduction of
Dioxygen to Hydrogen Peroxide by a Molecular Manganese Complex with a Bipyridine-
Containing Schiff Base Ligand. J. Am. Chem. Soc. 2018, 140, 3232-3241.
34.
Shook, R. L.; Peterson, S. M.; Greaves, J.; Moore, C.; Rheingold, A. L.; Borovik, A.
S., Catalytic Reduction of Dioxygen to Water with a Monomeric Manganese Complex at
Room Temperature. J. Am. Chem. Soc. 2011, 133, 5810-5817.
35.
Wasylenko, D. J.; Rodríguez, C.; Pegis, M. L.; Mayer, J. M., Direct Comparison of
Electrochemical and Spectrochemical Kinetics for Catalytic Oxygen Reduction. J. Am. Chem.
Soc. 2014, 136, 12544-12547.
36.
Asahi, M.; Yamazaki, S.-i.; Itoh, S.; Ioroi, T., Electrochemical reduction of dioxygen
by copper complexes with pyridylalkylamine ligands dissolved in aqueous buffer solution: the
relationship between activity and redox potential. Dalton Trans. 2014, 43, 10705-10709.
37.
Kopf, M. A.; Varech, D.; Tuchagues, J. P.; Mansuy, D.; Artaud, I., New intermediate-
spin chloroiron(III) complex with a mixed nitrogen-sulfur co-ordination. J. Chem. Soc.,
Dalton Trans. 1998, 991-998.
38.
Costas, M.; Mehn, M. P.; Jensen, M. P.; Que, L., Jr., Dioxygen Activation at
Mononuclear Nonheme Iron Active Sites: Enzymes, Models, and Intermediates. Chem. Rev.
2004, 104, 939-986.
39.
bond at iron complexes. Coord. Chem. Rev. 2017, 334, 2-24.
40. Hong, S.; Lee, Y.-M.; Ray, K.; Nam, W., Dioxygen activation chemistry by synthetic
Gamba, I.; Codola, Z.; Lloret-Fillol, J.; Costas, M., Making and breaking of the O-O
mononuclear nonheme iron, copper and chromium complexes. Coord. Chem. Rev. 2017, 334,
25-42.
41.
Ray, K.; Pfaff, F. F.; Wang, B.; Nam, W., Status of reactive non-heme metal-oxygen
intermediates in chemical and enzymatic reactions. J. Am. Chem. Soc. 2014, 136, 13942-
13958.
42.
Segaud, N.; Anxolabehere-Mallart, E.; Senechal-David, K.; Acosta-Rueda, L.; Robert,
M.; Banse, F., Electrochemical study of a nonheme Fe(II) complex in the presence of
dioxygen. Insights into the reductive activation of O2 at Fe(II) centers. Chem. Sci. 2015, 6,
639-647.
19
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 20 of 24
1
2
3
4
5
6
7
8
9
43.
Thibon, A.; England, J.; Martinho, M.; Young, V. G.; Frisch, J. R.; Guillot, R.; Girerd,
J.-J.; Muenck, E.; Que, L., Jr.; Banse, F., Proton- and reductant-assisted dioxygen activation
by a nonheme iron(II) complex to form an oxoiron(IV) intermediate. Angew. Chem., Int. Ed.
2008, 47, 7064-7067.
44.
Cavazza, C.; Bochot, C.; Rousselot-Pailley, P.; Carpentier, P.; Cherrier, M. V.; Martin,
L.; Marchi-Delapierre, C.; Fontecilla-Camps, J. C.; Menage, S., Crystallographic snapshots of
the reaction of aromatic C-H with O2 catalysed by a protein-bound iron complex. Nat. Chem.
2010, 2, 1069-1076.
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
45.
Brazzolotto, D.; CantúꢀReinhard, F. G.; Smith-Jones, J.; Retegan, M.; Amidani, L.;
Faponle, A. S.; Ray, K.; Philouze, C.; deꢀVisser, S. P.; Gennari, M.; Duboc, C., A High-
Valent Non-Heme μ-Oxo Manganese(IV) Dimer Generated from a Thiolate-Bound
Manganese(II) Complex and Dioxygen. Angew. Chem. Int. Ed. 2017, 56, 8211-8215.
46.
Addison, A. W.; Rao, T. N.; Reedijk, J.; Vanrijn, J.; Verschoor, G. C., Synthesis,
structure, and spectroscopic properties of copper(II) compounds containing nitrogen–sulphur
donor ligands; the crystal and molecular structure of aqua[1,7-bis(N-methylbenzimidazol-2′-
yl)-2,6-dithiaheptane]copper(II) perchlorate J. Chem. Soc., Dalton Trans. 1984, 1349-1356.
47.
Conroy-Lewis, F. M.; Simpson, S. J., Molecular mercaptan (thiol) complexes: the
synthesis of the tert-butanethiol complexes [Ru(η5-C5H5)(PPh3)(tert-BuNC)(tert-BuSH)]PF6
and [Ru(η5-C5H5)(dppm)(tert-BuSH)]PF6 and the x-ray single crystal structure of [Ru(η5-
C5H5)(dppm)(tert-BuSH)]PF6 [dppm = bis(diphenylphosphino)methane]. J. Chem. Soc.,
Chem. Commun. 1991, 388-9.
48.
Thiol Ligand. Inorg. Chem. 2002, 41, 2202-2208.
49. Ma, E. S. F.; Rettig, S. J.; Patrick, B. O.; James, B. R., Ruthenium(II) Thiol and H2S
McGuire, D. G.; Khan, M. A.; Ashby, M. T., Discontinuum between a Thiolate and a
Complexes: Synthesis, Characterization, and Thermodynamic Properties. Inorg. Chem. 2012,
51, 5427-5434.
50.
Allan, C. B.; Davidson, G.; Choudhury, S. B.; Gu, Z.; Bose, K.; Day, R. O.; Maroney,
M. J., Protonation and Alkylation of a Dinuclear Nickel Thiolate Complex. Inorg. Chem.
1998, 37, 4166-4167.
51.
Wang, L.; CantúꢀReinhard, F. G.; Philouze, C.; Demeshko, S.; deꢀVisser, S. P.; Meyer,
F.; Gennari, M.; Duboc, C., Solvent- and Halide-Induced (Inter)conversion between Iron(II)-
Disulfide and Iron(III)-Thiolate Complexes. Chem. Eur. J. 2018, 24, 11973-11982.
52.
Kovacs, J. A.; Brines, L. M., Understanding how the thiolate sulfur contributes to the
function of the non-heme iron enzyme superoxide reductase. Acc. Chem. Res. 2007, 40, 501-
509.
53.
Takamura, K.; Matsubara, C.; Matsumoto, T., Reaction specificity of a titanium(IV)-
porphyrin complex to hydrogen peroxide in view of an Ab initio study. Anal. Sci. 2008, 24,
401-404.
54.
Matsubara, C.; Kawamoto, N.; Takamura, K., Oxo[5, 10, 15, 20-tetra(4-
pyridyl)porphyrinato]titanium(IV): an ultra-high sensitivity spectrophotometric reagent for
hydrogen peroxide. Analyst 1992, 117, 1781-1784.
55.
Rodima, T.; Kaljurand, I.; Pihl, A.; Maeemets, V.; Leito, I.; Koppel, I. A., Acid-Base
Equilibria in Nonpolar Media. 2. Self-Consistent Basicity Scale in THF Solution Ranging
from 2-Methoxypyridine to EtP1(pyrr) Phosphazene. J. Org. Chem. 2002, 67, 1873-1881.
20
ACS Paragon Plus Environment

Page 21 of 24
Journal of the American Chemical Society
1
2
3
4
5
6
56.
Huan, T. N.; Simon, P.; Rousse, G.; Genois, I.; Artero, V.; Fontecave, M., Porous
dendritic copper: an electrocatalyst for highly selective CO2 reduction to formate in
water/ionic liquid electrolyte. Chem. Sci. 2017, 8, 742-747.
7
8
9
57.
Passard, G.; Ullman, A. M.; Brodsky, C. N.; Nocera, D. G., Oxygen Reduction
Catalysis at a Dicobalt Center: The Relationship of Faradaic Efficiency to Overpotential. J.
Am. Chem. Soc. 2016, 138, 2925-2928.
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
58.
Gerken, J. B.; Stahl, S. S., High-Potential Electrocatalytic O2 Reduction with
Nitroxyl/NOx Mediators: Implications for Fuel Cells and Aerobic Oxidation Catalysis. ACS
Cent. Sci. 2015, 1, 234-243.
59.
Maeda, Y., Isolation and crystal structure of a peroxo-bridged heme-copper complex. Angew.
Chem. Int. Ed. 2003, 42, 2788-2791.
60.
Chishiro, T.; Shimazaki, Y.; Tani, F.; Tachi, Y.; Naruta, Y.; Karasawa, S.; Hayami, S.;
Hematian, S.; Garcia-Bosch, I.; Karlin, K. D., Synthetic Heme/Copper Assemblies:
Toward an Understanding of Cytochrome c Oxidase Interactions with Dioxygen and Nitrogen
Oxides. Acc. Chem. Res. 2015, 48, 2462-2474.
61.
metalloproteins and insights from model complexes. J. Biol. Inorg. Chem 2017, 22, 253-288.
62. Zhang, X.; Furutachi, H.; Fujinami, S.; Nagatomo, S.; Maeda, Y.; Watanabe, Y.;
Kitagawa, T.; Suzuki, M., Structural and Spectroscopic Characterization of (μ-Hydroxo or μ-
Oxo)(μ-peroxo)diiron(III) Complexes:ꢁ Models for Peroxo Intermediates of Non-Heme Diiron
Proteins. J. Am. Chem. Soc. 2005, 127, 826-827.
Quist, D. A.; Diaz, D. E.; Liu, J. J.; Karlin, K. D., Activation of dioxygen by copper
63.
Yanhong, D.; Shiping, Y.; G., Y. V.; Lawrence, Q., Crystal Structure Analysis of a
Synthetic Non-Heme Diiron O2 Adduct: Insight into the Mechanism of Oxygen Activation.
Angew. Chem. Int. Ed. 1996, 35, 618-620.
64.
Y.; Ōkawa, H.; Hayashi, Y.; Uehara, A.; Suzuki, M., cis-μ-1,2-Peroxo Diiron Complex:ꢁ
Structure and Reversible Oxygenation. J. Am. Chem. Soc. 1996, 118, 701-702.
65.
porphyrins. Detection of a peroxo-bridged iron(III) porphyrin dimer and the mechanism of its
thermal decomposition to the oxo-bridged iron(III) porphyrin dimer. J. Am. Chem. Soc. 1980,
102, 4344-4350.
Ookubo, T.; Sugimoto, H.; Nagayama, T.; Masuda, H.; Sato, T.; Tanaka, K.; Maeda,
Chin, D.-H.; La Mar, G. N.; Balch, A. L., Mechanism of autoxidation of iron(II)
66.
Brines, L. M.; Shearer, J.; Fender, J. K.; Schweitzer, D.; Shoner, S. C.; Barnhart, D.;
Kaminsky, W.; Lovell, S.; Kovacs, J. A., Periodic trends within a series of five-coordinate
thiolate-ligated [MII(SMe2N4(tren))]+ (M = Mn, Fe, Co, Ni, Cu, Zn) complexes, including a
rare example of a stable CuII-thiolate. Inorg. Chem. 2007, 46, 9267-9277.
67.
Coggins, M. K.; Sun, X.; Kwak, Y.; Solomon, E. I.; Rybak-Akimova, E. V.; Kovacs,
J. A., Characterization of Metastable Intermediates Formed in the Reaction Between a Mn(II)
Complex and Dioxygen, Including a Crystallographic Structure of a Binuclear Mn(III)-
Peroxo Species. J. Am. Chem. Soc. 2013, 135, 5631-5640.
68.
Coggins, M. K.; Martin-Diaconescu, V.; DeBeer, S.; Kovacs, J. A., Correlation
Between Structural, Spectroscopic, and Reactivity Properties Within a Series of Structurally
Analogous Metastable Manganese(III) Alkylperoxo Complexes. J. Am. Chem. Soc. 2013,
135, 4260-4272.
21
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 22 of 24
1
2
3
4
5
6
69.
Coggins, M. K.; Kovacs, J. A., Structural and Spectroscopic Characterization of
Metastable Thiolate-Ligated Manganese(III) Alkylperoxo Species. J. Am. Chem. Soc. 2011,
133, 12470-12473.
7
8
9
70.
Lee, C.-M.; Chuo, C.-H.; Chen, C.-H.; Hu, C.-C.; Chiang, M.-H.; Tseng, Y.-J.; Hu,
C.-H.; Lee, G.-H., Structural and Spectroscopic Characterization of a Monomeric Side-On
Manganese(IV) Peroxo Complex. Angew. Chem. Int. Ed. 2012, 51, 5427-5430.
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
71.
Krishnamurthy, D.; Kasper, G. D.; Namuswe, F.; Kerber, W. D.; Narducci Sarjeant, A.
A.; Moenne-Loccoz, P.; Goldberg, D. P., A Low-Spin Alkylperoxo‚Iron(III) Complex with
Weak Fe-O and O-O Bonds: Implications for the Mechanism of Superoxide Reductase. J. Am.
Chem. Soc. 2006, 128, 14222-14223.
72.
Namuswe, F.; Kasper, G. D.; Sarjeant, A. A. N.; Hayashi, T.; Krest, C. M.; Green, M.
T.; Moenne-Loccoz, P.; Goldberg, D. P., Rational Tuning of the Thiolate Donor in Model
Complexes of Superoxide Reductase: Direct Evidence for a trans Influence in FeIII-OOR
Complexes. J. Am. Chem. Soc. 2008, 130, 14189-14200.
73.
Brown, C. D.; Neidig, M. L.; Neibergall, M. B.; Lipscomb, J. D.; Solomon, E. I.,
VTVH-MCD and DFT studies of thiolate bonding to {FeNO}(7)/{FeO2}(8) complexes of
isopenicillin N synthase: Substrate determination of oxidase versus oxygenase activity in
nonheme Fe enzymes. J. Am. Chem. Soc. 2007, 129, 7427-7438.
74.
Kitagawa, T.; Dey, A.; Lugo-Mas, P.; Benedict, J. B.; Kaminsky, W.; Solomon, E.;
Kovacs, J. A., A Functional Model for the Cysteinate-Ligated Non-Heme Iron Enzyme
Superoxide Reductase (SOR). J. Am. Chem. Soc. 2006, 128, 14448-14449.
75.
Compound II Is Basic: Implications for P450 Chemistry. Science 2004, 304, 1653-1656.
76. Sallmann, M.; Siewert, I.; Fohlmeister, L.; Limberg, C.; Knispel, C., A
Trispyrazolylborato Iron Cysteinato Complex as a Functional Model for the Cysteine
Dioxygenase. Angew. Chem. Int. Ed. 2012, 51, 2234-2237.
77.
H.; Limberg, C.; deꢀVisser, S. P., Structure and Mechanism Leading to Formation of the
Cysteine Sulfinate Product Complex of a Biomimetic Cysteine Dioxygenase Model. Chem.
Eur. J. 2015, 21, 7470-7479.
78.
thiolate complexes. Dalton Trans. 2012, 41, 10883-10899.
79. McQuilken, A. C.; Jiang, Y.; Siegler, M. A.; Goldberg, D. P., Addition of Dioxygen to
an N4S(thiolate) Iron(II) Cysteine Dioxygenase Model Gives a Structurally Characterized
Sulfinato–Iron(II) Complex. J. Am. Chem. Soc. 2012, 134, 8758-8761.
80.
Green, M. T.; Dawson, J. H.; Gray, H. B., Oxoiron(IV) in Chloroperoxidase
Sallmann, M.; Kumar, S.; Chernev, P.; Nehrkorn, J.; Schnegg, A.; Kumar, D.; Dau,
McQuilken, A. C.; Goldberg, D. P., Sulfur oxygenation in biomimetic non-heme iron-
McQuilken, A. C.; Matsumura, H.; Dürr, M.; Confer, A. M.; Sheckelton, J. P.; Siegler,
M. A.; McQueen, T. M.; Ivanović-Burmazović, I.; Moënne-Loccoz, P.; Goldberg, D. P.,
Photoinitiated Reactivity of a Thiolate-Ligated, Spin-Crossover Nonheme {FeNO}7 Complex
with Dioxygen. J. Am. Chem. Soc. 2016, 138, 3107-3117.
81.
Mediated Dioxygen Activation. Inorg. Chem. 2010, 49, 3646-3660.
82. Shook, R. L.; Borovik, A. S., The effects of hydrogen bonds on metal-mediated O2
Shook, R. L.; Borovik, A. S., Role of the Secondary Coordination Sphere in Metal-
activation and related processes. Chem. Commun. 2008, 6095-6107.
22
ACS Paragon Plus Environment

Page 23 of 24
Journal of the American Chemical Society
1
2
3
4
5
6
83.
Kobayashi, N.; Nevin, W. A., Electrocatalytic reduction of oxygen using water-soluble
iron and cobalt phthalocyanines and porphyrins. Appl. Organomet. Chem. 1996, 10, 579-590.
84. Soo, H. S.; Komor, A. C.; Iavarone, A. T.; Chang, C. J., A Hydrogen-Bond Facilitated
7
8
9
Cycle for Oxygen Reduction by an Acid- and Base-Compatible Iron Platform. Inorg. Chem.
2009, 48, 10024-10035.
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
23
ACS Paragon Plus Environment

Journal of the American Chemical Society
Page 24 of 24
1
2
3
4
Table of Contents
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
24
ACS Paragon Plus Environment