Mechanism of Catalytic O2 Reduction by Iron Tetraphenylporphyrin

Article abstract of DOI:10.1021/jacs.9b02640

The catalytic reduction of O₂ to H₂O is important for energy transduction in both synthetic and natural systems. Herein, we report a kinetic and thermochemical study of the oxygen reduction reaction (ORR) catalyzed by iron tetraphenylporphyrin (Fe(TPP)) in N,N′-dimethylformamide using decamethylferrocene as a soluble reductant and para-toluenesulfonic acid (pTsOH) as the proton source. This work identifies and characterizes catalytic intermediates and their thermochemistry, providing a detailed mechanistic understanding of the system. Specifically, reduction of the ferric porphyrin, [Fe^III(TPP)]+ , forms the ferrous porphyrin, Fe^II(TPP), which binds O₂ reversibly to form the ferricsuperoxide porphyrin complex, Fe^III(TPP)(O₂ ?-). The temperature dependence of both the electron transfer and O₂ binding equilibrium constants has been determined. Kinetic studies over a range of concentrations and temperatures show that the catalyst resting state changes during the course of each catalytic run, necessitating the use of global kinetic modeling to extract rate constants and kinetic barriers. The rate-determining step in oxygen reduction is the protonation of Fe^III(TPP)(O₂ ?-) by pTsOH, which proceeds with a substantial kinetic barrier. Computational studies indicate that this barrier for proton transfer arises from an unfavorable preassociation of the proton donor with the superoxide adduct and a transition state that requires significant desolvation of the proton donor. Together, these results are the first example of oxygen reduction by iron tetraphenylporphyrin where the pre-equilibria among ferric, ferrous, and ferric-superoxide intermediates have been quantified under catalytic conditions. This work gives a generalizable model for the mechanism of iron porphyrin-catalyzed ORR and provides an unusually complete mechanistic study of an ORR reaction. More broadly, this study also highlights the kinetic challenges for proton transfer to catalytic intermediates in organic media.

Full text of DOI:10.1021/jacs.9b02640

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Article
The Mechanism of Catalytic O2 Reduction by Iron Tetraphenylporphyrin
Michael L. Pegis, Daniel J. Martin, Catherine F. Wise, Anna C. Brezny, Samantha I.
Johnson, Lewis E. Johnson, Neeraj Kumar, Simone Raugei, and James M. Mayer
J. Am. Chem. Soc., Just Accepted Manuscript • Publication Date (Web): 01 May 2019
Downloaded from http://pubs.acs.org on May 1, 2019
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The Mechanism of Catalytic O2 Reduction by
Iron Tetraphenylporphyrin
Michael L. Pegis,1,3,§ Daniel J. Martin,1,§ Catherine F. Wise,1,§ Anna C. Brezny,1 Samantha I.
Johnson,2‡ Lewis E. Johnson,4 Neeraj Kumar,5 Simone Raugei,2 and James M. Mayer1*
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Abstract
The catalytic reduction of O2 to H2O is important for energy transduction in both synthetic and
natural systems. Herein, we report a kinetic and thermochemical study of O2 reduction catalyzed
by iron tetraphenylporphyrin (Fe(TPP)) in N,N′-dimethylformamide using decamethylferrocene
as a soluble reductant and para-toluenesulfonic acid (pTsOH) as the proton source. This work
identifies and characterizes catalytic intermediates and their thermochemistry, providing a
detailed mechanistic understanding of the system. Specifically, reduction of the ferric porphyrin,
[FeIII(TPP)]+ forms the ferrous porphyrin, FeII(TPP), which binds O2 reversibly to form the ferric-
,
•‒
superoxide porphyrin complex, FeIII(TPP)(O2 ). The temperature dependence of both the
electron transfer and O2 binding equilibrium constants have been determined. Kinetic studies
over a range of concentrations and temperatures show that the catalyst resting state changes
during the course of each catalytic run, necessitating the use of global kinetic modeling to extract
rate constants and kinetic barriers. The rate-determining step in oxygen reduction is protonation
•‒
of FeIII(TPP)(O2 ) by pTsOH, which proceeds with a substantial kinetic barrier. Computational
studies indicate that this barrier for proton transfer arises from an unfavorable pre-association
of the proton donor with the superoxide adduct and a transition state that requires significant
desolvation of the proton donor. Together, these results are the first example of oxygen
reduction by iron tetraphenylporphyrin where the pre-equilibria between ferric, ferrous, and
ferric-superoxide intermediates have been quantified under catalytic conditions. This work gives
a generalizable model for the mechanism of iron porphyrin-catalyzed ORR and provides an
unusually complete mechanistic study of an ORR reaction. More broadly, this study also
highlights the kinetic challenges for proton transfer to catalytic intermediates in organic media.
Introduction
Many diverse biological and energy processes involve the catalytic reduction of dioxygen
(O2).1–5 In nature, cytochrome c oxidase reduces O2 to drive ATP synthesis in cellular respiration,6
and cytochromes P450 couple the reduction of O2 to the oxidations of endogenous and
xenobiotic substances in numerous synthetic and metabolic processes.7–9 In energy conversion
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technologies such as fuel cells, the oxidation of a fuel (dihydrogen, methanol, etc.) is coupled to
the oxygen reduction reaction (ORR, eq 1), producing usable electrochemical work for portable
and stationary applications.10 For such approaches to be practical, the ORR must be performed
at fast rates and with high selectivity and energy efficiency over thousands of hours of operation.
Achieving this goal will have vast implications for the global energy economy.
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O2 + 4H + + 4푒 ‒ ⟶2H2O
(1)
Given the kinetic complexity of the ORR, widespread commercialization of fuel cell
technologies requires identification of inexpensive and efficient electrocatalysts capable of
delivering 4H+ and 4e‒ to O2 and cleaving the O=O bond. Numerous research efforts have focused
on developing new catalytic systems and improving previously known ORR
electrocatalysts.3,4,11,12 Given their biological relevance as active sites, iron porphyrins have been
widely studied as ORR catalysts under electrocatalytic conditions.13,14 Notably, graphite-
embedded iron porphyrinic-like materials have been shown to exhibit ORR activity rivaling that
of platinum metal in acidic media; however, the exact nature of their active sites remains a
continued discussion.6-8
Further improvement of ORR cathodic materials will require an understanding of the
catalyst identity and turnover-limiting step(s) in the catalytic cycle. Such a detailed understanding
is difficult to achieve for heterogeneous electrode surfaces, where the nature and catalytic
activities of specific sites and intermediates are difficult to determine. Studying molecular ORR
electrocatalysts that have been physi- or chemisorbed onto electrode surfaces can circumvent
some of these complications, especially with the use of in-situ spectroscopic characterization.18–
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Reaction mechanisms are more easily analyzed in systems with homogeneous molecular
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catalysts, as these catalysts can be prepared in pure form and are readily examined using
standard spectroscopic and kinetic methods for solution species. Optical spectroscopies have
been extensively used to study iron porphyrins in the biomimetic context of dioxygen binding24–
26 and the oxidation of organic molecules.27,28 However, surprisingly few studies have investigated
the catalytic reduction of O2 by iron porphyrins using soluble reductants.11
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Nearly 30 years ago, two reports by Fukuzumi described the first mechanistic
investigations of O2 reduction by iron tetraphenylporphyrin (abbreviated herein as Fe(TPP) if the
oxidation state and axial ligand(s) are not specified). These studies were performed in acetonitrile
(MeCN) using substituted ferrocenes (Fc) as the source of electrons and perchloric acid (HClO4)
as the source of protons.29,30 Under these conditions, the rate-determining step was proposed to
be outer-sphere electron transfer from the reductant to [FeIII(TPP)]ClO4, based upon the first-
order dependence on [Fc] and the zero-order dependences on [O2] and [HClO4]. In this example,
the strongly acidic medium with relatively weak reductants resulted in no detectable
intermediates within the catalytic cycle.
More recently, kinetic studies examined O2 reduction catalyzed by FeIII(TPP)Cl and other
porphyrin derivatives in acidic acetonitrile and N,N′-dimethylformamide solutions containing
decamethylferrocene (Fc*) as a soluble reductant.31,32 With excess of this stronger reductant, the
rate law for catalysis via FeIII(TPP)Cl was zero-order in reductant because electron transfer from
Fc* to FeIII(TPP)Cl was initially rapid and favorable. Under these conditions, the rate of ORR
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catalysis was first order in O2, acid, and catalyst.31 The same rate law and rate constant were
obtained under electrocatalytic conditions. Subsequent studies found this rate law to be general
for electrocatalytic O2 reduction using substituted iron porphyrins32 and in the presence of
proton donors of varying acidity (pKa).33 The rate law implicates a mechanism of initial reduction
of FeIII to FeII, pre-equilibrium O2 binding to FeII, and rate-limiting proton transfer from the acid
to the FeIII-superoxo intermediate (Scheme 1). However, efforts to decrease the proton transfer
barrier by appending proton relays to the iron porphyrins did not result in rate enhancements.32
To obtain direct support for the proposed mechanism and to better understand the factors
affecting catalytic rates, a more thorough mechanistic analysis was needed.
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Scheme 1. Proposed mechanism for oxygen reduction catalyzed by Fe(TPP), with TPP
abbreviated as an oval. Red = Fc* or electrode.
3Red+, 3pTsO , 2H2O
3Red, 3pTsOH
FeIII
Red
Fast
Red+
OH
O
FeIII
FeII
Rate-Determining
PT
KO2
kPT
O2
O
pTsO
O
FeIII
Pre-equilibrium
O2 Binding
pTsOH
Herein, we report a much more complete and nuanced analysis of the 4H+/4e– reduction
of O2 to H2O catalyzed by iron tetraphenylporphyrin, prepared as the triflate salt ([FeIII(TPP)]OTf).
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The ferrous and ferric-superoxide porphyrin species are identified and observed as catalytic
intermediates for the first time, and their temperature-dependent speciation is reported. This
includes measurements of the equilibrium constants for electron transfer from Fc* to
[FeIII(TPP)]OTf and for O2 binding to FeII(TPP). These shifting equilibria and speciation play a
critical and previously unappreciated role in the kinetics of catalysis, as the catalyst resting state
varies among three different species during the course of the reaction. The temperature-
dependence of these equilibria combined with measurements of catalytic rates and catalyst
speciation by optical stopped-flow and electrochemical methods provide an unusually rich
dataset. Modelling these results gives a detailed view of the catalytic mechanism under different
conditions. The model gives the activation parameters for protonation of the ferric-superoxo
intermediate, the rate-determining step in the catalytic cycle. These results are used in
conjunction with computational modeling to analyze the underlying factors limiting proton
transfer to the ferric-superoxide intermediate. Such a detailed mechanistic understanding
provides critical insight into iron porphyrin-catalyzed oxygen reduction and should enable further
improvement of catalytic systems.
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O2 + 4푝TsOH + 4Fc ⋆ [Fe(TPP)]OTf2H2O +4푝TsO ‒ + 4Fc ⋆ +
(2)
Experimental
Materials
Iron tetraphenylporphyrin chloride, FeIII(TPP)Cl, was synthesized by the metalation of
meso-tetraphenylporphyrin with iron(II) chloride in refluxing DMF as previously reported32 (see
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SI, Figure S1-S2). [FeIII(TPP)]OTf was generated in situ from Fe(TPP)Cl and TlOTf in acidified DMF
(see SI, Figure S3-S4). FeII(TPP) was prepared from the well-known reduction of FeIII(TPP)Cl by
Zn/Hg amalgam and was recrystallized from toluene/THF mixtures.34,35 This compound was also
prepared using modified literature procedures (see SI, Figure S7-S8). All solvents and electrolytes
were purchased at high purity and purified as needed, as further described in Section 1.2 of the
SI.
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Optical Equilibrium and Kinetic Measurements
Equilibrium measurements were performed by optical spectroscopy using a temperature-
controlled Unisoku Unispeks cryostat. Equilibrium constants for electron transfer were made by
fitting the UV-vis spectra of DMF solutions containing 0.1 M [n-Bu4N][PF6] and equimolar (100
μM) amounts of [FeIII(TPP)]OTf, FeII(TPP), Fc*, and Fc*+ between 213 K and 293 K. Likewise,
equilibrium constants for O2 binding were made by fitting the UV-vis spectra of O2-saturated DMF
solutions initially containing 0.1 M [n-Bu4N][PF6] and 60 μM FeII(TPP) between 213 K and 238 K.
Both sets of equilibria data were analyzed using van ‘t Hoff plots to yield enthalpic and entropic
parameters. See Sections 2-3 of the SI for complete details.
Stopped-flow kinetics measurements were performed by mixing N2-saturated DMF
solutions of Fc* in one syringe with air-saturated DMF solutions containing [FeIII(TPP)]OTf and
pTsOH in a second syringe. The reaction progress was monitored by following the appearance of
Fc*+ over time. Experiments were performed at various [pTsOH], [catalyst], and temperature (253
– 303 K). Changes in catalyst speciation over time were determined by fitting optical spectra to
•‒
linear combinations of three catalyst species: [FeIII(TPP)]OTf, FeII(TPP), and FeIII(TPP)(O2 ). The
concentration time courses were modeled to obtain thermodynamic and kinetic parameters. See
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Section 5-6 of the SI for complete details.
Kinetic Modeling
The time course data from the stopped-flow catalytic experiments were fit to a multi-step
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kinetic model in the program COPASI.36 The thermodynamic parameters (∆Ho and ∆So) for
electron transfer and O2 binding, as well as the activation parameters (∆H‡ and ∆S‡) for proton
transfer, were optimized in order to minimize the sum of squared deviations between the
experimental concentrations and fitted time course concentrations. See Section 6 of the SI for
complete details.
Computational Modeling
The present density functional theory (DFT) calculations build on previously reported
work.32 Stationary points were optimized using the PBE exchange and correlation functional.37
The Stuttgart/Dresden basis set with relativistic effective core potential (SDD) was used for the
Fe center, and the 6-31G** basis set38,39 was used for all atoms except sulfur. Diffuse functions
and additional polarization functions (6-31++G(2df,2pd)) were used on sulfur in order to capture
the sulfur–oxygen bonding correctly. Justification for this can be found in Section 8 of the SI.
Single point solvation energies in DMF were modeled using the SMD continuum solvent.40
Harmonic vibrational frequencies, calculated at the same level of theory, were used to estimate
zero-point energy (ZPE) and the thermal contributions free energies. Free energies are referred
to the standard state concentration of 1 M for the solute and 12.9 M for the solvent DMF at T =
298 K. See Section 8 of the SI for complete details.
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Results
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I. Electron Transfer Equilibrium
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The proposed mechanism involves initial reduction of [FeIII(TPP)]OTf to FeII(TPP) by
decamethylferrocene (Fc*). Qualitatively, addition of excess Fc* to [FeIII(TPP)]OTf results in the
predominant formation of FeII(TPP) (Figure S9-S10). Adding Fc*+ to this solution reoxidizes some
FeII(TPP) and confirms the equilibrium nature of this ET process (eq 3; Figure S11). Equilibrium
constants in DMF were measured by preparing solutions containing the four species and 0.1 M
[n-Bu4N][PF6], to make the conditions similar to those used in the electrocatalytic and optical
measurements. After correcting for dilution, analysis of the optical spectra gave the ratio of
[FeIII(TPP)]OTf to FeII(TPP) (see SI).
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Fc ⋆ + FeIII(TPP) + ⇌Fc ⋆ + + FeII(TPP)
(3)
∗ +
II
[
Fc
]
[Fe (TPP)]
(4)
퐾ET = [Fc ∗ ][[FeIII(TPP)] +
]
Following eq 4 and knowing the concentrations of all four species,
was determined
퐾ET
at 293 K to be 0.16 ± 0.03. This
could also be determined from electrochemical
퐾ET
measurements: E1/2 (FeIII(TPP)+/FeII(TPP)) = ‒0.538 V vs. Fc+/Fc (vide infra) and E1/2(Fc*+/Fc*) =
‒0.484 V vs. Fc+/Fc in DMF31 containing 0.1 M [n-Bu4N][PF6]. The 0.054 V difference in reduction
potentials gives
= 0.12 (SI Section 2.2), very close to the
= 0.16 from optical
퐾ET
퐾ET
measurements. The optical measurements were repeated at temperatures between 213 and 293
K to investigate the temperature dependence of (Figure 1a). The resulting van ‘t Hoff plot
퐾ET
yielded the enthalpy and entropy for the electron transfer from Fc* to [FeIII(TPP)]OTf: ΔH°ET = 2.8
± 0.1 kcal mol-1 and ΔS°ET = 6 ± 2 cal mol-1 K-1 (Figure 1b).
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Figure 1. a) Spectra of a DMF solution containing 100 μM [FeIII(TPP)]OTf, FeII(TPP), Fc*, and Fc*+
and 0.1 M [n-Bu4N][PF6] in the Q-band region at temperatures between 293 K and 213 K. b) van
‘t Hoff analysis of ln(
) derived from the data in a).
퐾ET
II. Oxygen Binding Equilibrium
To test the hypothesis that O2 binding follows the initial reduction of [FeIII(TPP)]OTf, a
genuine FeII(TPP) sample was prepared and exposed to O2 at low temperatures in DMF. The
addition of 1 atm O2 to DMF solutions of FeII(TPP) at 213 K resulted in the quantitative conversion
of FeII(TPP) to a new species with Q-band absorbance features at 542 and 580 nm, characteristic
•‒
of a ferric tetra-arylporphyrin superoxide complex, FeIII(TPP)(O2 ).41,42 The chemically-generated
FeIII(TPP)(O2 ) complex was relatively stable (~30 minutes) at 213 K but decomposed quickly at
•‒
temperatures above 238 K to the well-known μ-oxo dimer, [FeIII(TPP)]2O.42
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The stoichiometry of O2 binding was measured by titrating a DMF solution of FeII(TPP) at
213 K with a room temperature DMF solution containing 1 atm O2 (3.1 mM [O2]).43 With
substoichiometric concentrations of O2, the Q-bands of FeII(TPP) incrementally shifted to those
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•‒
of FeIII(TPP)(O2 ) (Figure 2a). At 213 K, slightly more than one equivalent of O2 was necessary to
reach the end point of the titration, after which no additional changes occurred in the spectra
(Figure S13 and S15).
The reversibility of O2 binding to FeII(TPP) was investigated by sparging a solution of the
•‒
generated FeIII(TPP)(O2 ) adduct with argon at 213 K. A spectrum collected after 5 minutes of
sparging closely matched the initial anaerobic spectrum of FeII(TPP) with some residual
contributions from remaining FeIII(TPP)(O2 ) (Figure S16). That FeII(TPP) could be formed from
•‒
FeIII(TPP)(O2 ) proved O2 binding reversibility and enabled us to measure equilibrium constants
•‒
(eq 5).
‒
FeIII(TPP)
(
O2
)
]
•
[
―1
(5)
(
M
)
=
퐾O
2
II
[ ]
Fe (TPP) O2
[
]
The equilibrium constants for O binding to FeII(TPP) ( , eq 5) were measured at various
퐾O
2
2
temperatures. With excess O2, a solution of FeIII(TPP)(O2 ) was prepared at 213 K and
•‒
incrementally warmed to 238 K. Upon warming, the initial spectrum gradually changed to contain
Q-bands of both FeIII(TPP)(O2 ) and FeII(TPP) (Figure S17). The UV-vis spectrum at each
•‒
temperature was fit to a linear combination of independently measured FeIII(TPP)(O2 ) and
•‒
FeII(TPP) spectra (see SI Section 3.3). The concentrations of FeII(TPP), FeIII(TPP)(O2 ), and of
•‒
dissolved O from mass balance were then used to calculate
at each temperature (eq 5). At
퐾
2
O2
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•‒
213 K, the free energy of formation for FeIII(TPP)(O2 ) is exergonic, ΔGo 2 = -3.68 kcal mol-1. The
O
5
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slope and intercept of the resulting van ‘t Hoff plot (Figure 2b) yielded the enthalpy and entropy
•‒
for the binding of O to form FeIII(TPP)(O ): ΔH° = –10.5 ± 0.7 kcal mol-1 and ΔS° = –32 ± 3 cal
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2
2
O
2
2
O
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mol–1 K–1. By extrapolating the van ’t Hoff plot to 298 K , the equilibrium constant for O2 binding
at room temperature was determined to be ca. 5 M–1. A T-test gave a 95% confidence interval
for this extrapolated value at 298 K to be between 0.21 M-1 and 110 M-1 (see SI for details).44
Figure 2. a) Optical spectra (corrected for dilution) of a titration at 213 K of a DMF solution
containing 50 μM FeII(TPP) and 0.1 M [n-Bu4N][PF6] with DMF containing dissolved O2. b) van ‘t
Hoff analysis of O binding to FeII(TPP) (
in M–1).
퐾
2
O2
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III. Electrochemical Kinetics of Catalytic Oxygen Reduction
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Our previous studies have shown that the rate of ORR catalyzed by iron tetra-
arylporphyrins exhibits first-order dependences on the concentrations of catalyst, O2, and acid.32
In order to verify that the same rate law applies to the system studied here, we investigated the
kinetics of ORR catalyzed by [FeIII(TPP)]OTf using cyclic voltammetry.
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In anaerobic DMF solutions containing [FeIII(TPP)]OTf and 0.1 M [n-Bu4N][PF6], a
chemically reversible, diffusion-controlled reduction was observed at –0.538 V vs. Fc+/Fc (Figure
3a (black trace), Figure S20). This wave is diagnostic of the reduction of [FeIII(TPP)]+ to FeII(TPP),
as previously reported by Savéant45 and later by our group.32,33 Addition of pTsOH did not affect
E1/2 (FeIII(TPP)+/FeII(TPP)) (Figure S21), in contrast to what has been observed for Fe(TPP) in the
presence of chloride salts.31 These results are consistent the preferred binding of DMF solvent
over triflate, in both Fe(III) and Fe(II) redox states, as previously observed.32,45 For simplicity, we
exclude explicit solvent molecules from the abbreviated formulas and continue to abbreviate
these redox states as [FeIII(TPP)]OTf and FeII(TPP).
When DMF solutions containing [FeIII(TPP)]OTf, pTsOH, and [n-Bu4N][PF6] were sparged
with 1 atm air (0.7 mM O2 in DMF),43 the voltammograms displayed significant current
enhancements (>10x) centered above E1/2(FeIII/FeII) (Figure 3a, red trace). The magnitude of this
irreversible current was linear with the square root of the scan rate, which suggests that the
catalytic current arises from an electrogenerated, homogeneous catalyst.31 Rotating ring-disk
electrochemistry (RRDE) was used to quantify the amount of H2O2 produced during turnover. The
RRDE measurements demonstrated that catalysis is selective for the 4H+/4e– reduction of O2 to
water and produces ~1.5% H2O2 (ncat = 3.9) under these conditions (Figure S50-S51).
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Figure 3. a) Electrochemical response of 0.3 mM [FeIII(TPP)]OTf in the presence of 1 M pTsOH and
1.6 mM O2 (green), 0.7 mM O2 (red, air sparged), and 0 mM O2 (black, N2 sparge) at 298 K. b)
Dependence of kobs from FOWA on [pTsOH] at different O2 concentrations. All CVs were taken in
the presence of 0.1 M [n-Bu4N][PF6] electrolyte.
The pseudo-first order rate constant kobs for ORR was quantified using foot-of-the-wave
analysis (FOWA, equation 7),46 where ncat = 3.9 (determined by RRDE, see above). FOWA allows
the determination of catalytic rate constants under conditions where the concentrations of the
substrates in the reaction-diffusion layer (O2 and pTsOH in this case) are essentially identical to
the bulk solution concentrations (see SI).46 The linearity of the foot-of-the-wave plots (Figure S22)
indicated that the reaction was first order in catalyst. Additionally, kobs was linearly dependent
on both [O2] and [pTsOH] (Figure 3), indicating the third-order rate law in equations 6 and 8. The
quotient of kobs and the reactant concentrations then afforded the third order rate constant kcat.
Considering the upper and lower bounds for  in eq 7 (see SI), kcat ranged from (3.2-6.4)  105
M-2 s-1.
푟푎푡푒 = 푛cat푘obs[Fe(TPP)]
(6)
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푅푇
퐹푣
휎
2.24푛cat
푘obs
푖푐
(7)
(8)
=
푖푝
퐹
푅푇
(
)
]
1 + exp [
퐸 ― 퐸1/2
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[ ]
푘obs = 푘cat O2 [푝TsOH]
IV. Optical Kinetics
The kinetics of Fe(TPP)-catalyzed ORR were also studied by optical spectroscopy, over a
range of temperatures. These experiments were performed in DMF using pTsOH as the proton
source and Fc* as the soluble reductant. During optical measurements, all solutions also
contained 0.1 M [n-Bu4N][PF6] to mirror the electrochemical conditions. In a typical experiment,
a solution of air-saturated DMF (0.7 mM O2) containing [FeIII(TPP)]OTf (60 μM) and pTsOH (50
mM) was rapidly mixed in a stopped-flow instrument with an equal volume of an anaerobic DMF
solution of Fc* (6 mM, higher concentrations were limited by solubility), with all concentrations
being halved upon mixing (see SI). After mixing, optical spectra showed the growth of a broad
absorbance feature between 600-700 nm, diagnostic of decamethylferrocenium (Fc*+)
formation, along with various changes in the porphyrin Q-band region (Figures 4a and S24,
discussed further below). Control experiments performed under identical conditions but in the
absence of [FeIII(TPP)]OTf catalyst formed a much smaller amount of Fc*+ over the same time
course (~10%) (Figure 4b). The small amount of Fc* auto-oxidation likely results from unfavorable
•‒ 47
pre-equilibrium electron transfer with O2 followed by irreversible chemical reactions of O2 .
When pTsOD was substituted for pTsOH, the rate of Fc*+ formation did not change (KIE ≈ 1, see
Figure S25).
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Figure 4. a) Spectral changes measured for the reaction of O2 (0.33 mM) with pTsOH (50 mM)
and Fc* (3 mM), catalyzed by [FeIII(TPP)]OTf (30 M) over a 30 second reaction at 298 K. b)
Rate of Fc*+ formation over the course of the first 10 s of reaction (a) (black trace), compared to
the same reaction conducted in the absence of added [FeIII(TPP)]OTf catalyst (purple trace).
Concentrations of Fc*+ were calculated at 700 nm after removal of the absorbance
contributions from [FeIII(TPP)]OTf (see SI, Section 5.3).
The kinetics of the Fe(TPP)-catalyzed ORR were examined by stopped-flow methods at
temperatures between 253 and 303 K. In these experiments, the air-saturated syringe was
prepared at room temperature so that the initial concentration of dissolved O2 remained
constant for all catalytic runs. Notably, as shown in Figure 5, the spectrum collected 1 s after
mixing was temperature dependent. At 253 K, absorbance features at 542 and 580 nm were
immediately present (Figure 5a). These features match those of the independently prepared
•‒
ferric superoxide, FeIII(TPP)(O2 ), described above. However, at temperatures above 273 K, the
•‒
initial spectrum resembled a mixture of FeII(TPP) and FeIII(TPP)(O2 ) (Figure 5b). At all
temperatures, the initial traces decayed with the formation of [FeIII(TPP)]OTf and Fc*+.
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Figure 5. a) Spectral changes after stopped-flow mixing of [FeIII(TPP)]OTf (30 μM), pTsOH (35
mM), Fc* (3 mM), and O2 (0.33 mM) at 253 K. Black trace = initial spectra after mixing (1 s) Red
trace = final spectra after mixing (50 s). b) Same conditions as (a), at 303 K. Black trace = initial
spectra after mixing (1 s) Red trace = final spectra after mixing (20 s).
The temperature dependence of the initial stopped-flow kinetic traces reflects the
temperature-dependent equilibrium constant for O2 binding. The initial spectra showed
•‒
•‒
exclusively FeIII(TPP)(O2 ) at low temperatures and a mixture of FeII(TPP) and FeIII(TPP)(O2 ) at
higher temperatures, consistent with the independent O2 binding measurements described
above. While the electron transfer equilibrium also shows a temperature dependence, GET only
changes by ~ 0.3 kcal mol-1 across this temperature range. Thus, the observed spectral differences
are dominated by the difference in O2-binding favorability between the two temperatures.
•‒
Additionally, the observation of FeII(TPP) or FeIII(TPP)(O2 ) as the initial catalyst resting state also
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supports the proposed mechanism in which the rate of ORR is limited by protonation of
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7
•‒
FeIII(TPP)(O2 ) rather than initial electron transfer or O2 binding (Scheme 1).
8
9
Over the course of the stopped-flow kinetics runs, the speciation of the iron catalyst
changed, with [FeIII(TPP)]OTf becoming the predominant species at the end of the reaction (when
limiting O2 is completely consumed). The changes in catalyst speciation reflect the positions of
the time- and temperature-dependent equilibria between [FeIII(TPP)]OTf, FeII(TPP), and
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•‒
FeIII(TPP)(O2 ) during catalytic turnover, as discussed in more detail in the next section. In the
absence of a constant catalyst resting state, the rate law for Fc*+ formation needs to be defined
with respect to total amount of catalyst in solution, [Fe(TPP)]total. The resulting rate law (eq 9)
includes the two pre-equilibria (
퐾ET
and
) and the turnover-limiting protonation step, kPT. The
퐾O
2
complexity of this rate law arises from the uphill electron transfer equilibrium between Fc* and
[FeIII(TPP)]OTf (
< 1). As such, even with a large excess of Fc* (e.g. >40:1 ratio of [Fc*]:[O2]),
퐾ET
the growth of the oxidized product Fc*+ shifts the initial electron transfer reaction away from the
formation of catalytically active FeII(TPP), as demonstrated independently in Section I above. The
complexity of this rate law and the inability to find experimental conditions under which it
simplifies have precluded simple fitting and the determination of a simple reaction order in [Fc*]
(see SI).
⋆
푑[Fc ⋆ +
]
[ ]
[
]
푛cat퐾ET퐾O 푘PT[Fe(TPP)]total O2 Fc [푝TsOH]
2
(9)
=
⋆
⋆ +
푑푡
[ ]
퐾ET Fc (1 + 퐾O O2 ) + Fc
[
]
[
]
2
Longer time-scale catalytic experiments were conducted by combining a large excess of
Fc* (10 mM) and pTsOH (260 mM) relative to [FeIII(TPP)]OTf (1 M) and stirring the solution in
ambient air. Over 15 minutes, significantly larger amounts of Fc*+ were produced than an
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identical experiment conducted in the absence of [FeIII(TPP)]OTf. Iodometric titrations of the
resulting solution were used to quantify the amount of H2O2 produced during the catalyzed
reaction and revealed that ~15% H2O2 (ncat = 3.7 e–/O2, Figure S51) was formed under such
conditions, in good agreement with the selectivity values obtained electrochemically (ncat ≈ 3.9
e–/O2).31 These results demonstrate that the catalyst remains active for ORR for at least 15
minutes and permit a rough estimation of the catalyst turnover number over 15 minutes, TON =
2000 moles O2 consumed per mole catalyst, Figure S26.
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V. Kinetic Modeling
The constantly evolving catalyst speciation during ORR by Fe(TPP), observed by stopped-
flow measurements, required us to use a complete kinetic model to fit the data. As described in
this section, the model revealed the thermodynamic parameters for the ET and O2-binding pre-
equilibria and the kinetic parameters for the turnover-limiting step.
Fitting the optical data to a kinetic model first required the concentration of each catalyst
species at every time point to be determined. These concentrations were obtained using Beer’s
Law and a system of linear equations that considered absorbance contributions from
•‒
[FeIII(TPP)]OTf, FeII(TPP), FeIII(TPP)(O2 ), and Fc*+. Absorbance contributions from Fc* were
negligible in the wavelength region of interest and could be ignored (see SI Section 5.3).
Global modeling of all the room temperature stopped-flow kinetic runs was performed
with the software COPASI36 using the kinetic model in Scheme 2 and the corresponding rate law
(eq 9 above). The only parameters input into the model were the fast rate constants (107 M-1 s-1)
for (1) electron transfer from Fc* to FeIII(TPP)+ and (2) O2-binding to FeII(TPP). This was done to
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ensure that these steps were fast pre-equilibria, as observed experimentally (see above). Step 4
was included to account for mass balance in the reaction and was set to a fast enough rate to be
kinetically invisible.
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Scheme 2. Kinetic model and parameters used for global fitting of the stopped-flow kinetic
data, to and obtain ∆H°ET, ∆S°ET, ∆H°O2, ∆S°O2, ∆H‡ , and ∆S‡ .
PT
PT
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Table 1. Results of COPASI global fitting of stopped flow optical kinetic results and comparison
with experimental values.a
6
7
8
9
Reaction
Parameter b
(298 K)
Model
Experimental c
0.055(7)
2.9(1)
0.18(3)
2.8(1)
6(2)
퐾ET
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Fc* + FeIII(TPP)+ ⇌
∆H°ET
∆S°ET
Fc*+ + FeII(TPP)
4.1(4)
(250 K, M-1)d
(0.205 - 11.5)  102
퐾
O2
(2.7 - 6.4)  104
e
FeII(TPP) + O2 ⇌
∆H°O2
∆S°O2
–9.5(1)
–10.5(7)e
–32 (3)e
III
–
•
Fe (TPP)(O2 )
–16.9(4)
1.5(3)
2.0(9)f
kPT (298 K)
 103 M-1 s-1
 105 M-1 s-1
III
–
•
Fe (TPP)(O2 ) + pTsOH ⟶
∆H‡
12.0(1)
-3.1(2)
–
–
III
+
–
PT
•
Fe (TPP)(O2H ) + pTsO
∆S‡
PT
FeII(TPP) + O2 + pTsOH ⟶
kcat (298 K, M-2 s-1)
kPT
(2-20)  105 h
(3.2‒6.4)  105 i
=
퐾O
III
+
– g
•
2
Fe (TPP)(O2H ) + pTsO
a
Optimized values from COPASI analysis of data at 253-303 K, 50-100 mM pTsOH, 30–50 μM
Fe(TPP). COPASI values and experimental equilibrium parameters are given with the uncertainty
in parentheses representing one standard deviation. ∆H° and ∆H‡ values in kcal mol-1; ∆S° and
b
∆S‡ values in cal K–1 mol-1. Direct experimental measurements of equilibrium parameters from
c
optical
spectra
and
van
‘t
Hoff
analyses
(Sections
I
and
II).
d
Values extrapolated from thermodynamic parameters from kinetic data (Model) or
e
lower-temperature equilibrium measurements (Experimental). The 95% confidence limits from
a T-test are for 2 at 298 K: 0.21 to 110 M–1; for ∆H°O2: -9.6 to -11.8 kcal mol-1; and for ∆SO2: -28
퐾O
to -38 cal K-1 mol-1, SI Section 3.4). f Calculated as kcat(echem)
/
, where
was defined
퐾
O2
퐾O
2(experimental)
as the mean value, 5 M-1. Chemical steps involved in defining the catalytic rate constant, kcat,
g
h
i
determined electrochemically. Calculated as
of-the-wave analysis, vide infra.
(model)  kPT (model). Calculated from foot-
퐾O
2
The model used in Scheme 2 attempted to optimize both the simulated rate of
decamethylferrocenium formation (d[Fc*+]/dt) and the catalyst speciation to the experimental
data by varying the thermodynamic parameters for electron transfer (∆H°ET and ∆S°ET, step 1),
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the thermodynamic parameters for O2 binding ((∆H°O2 and ∆S°O2, step 2) and the activation
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parameters for proton transfer (∆H‡PT and ∆S‡ , step 3). The six parameters were optimized to
PT
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fit all stopped flow data simultaneously—including experiments between 253 and 303 K with
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varied concentrations of substrate (50–100 mM pTsOH) and catalyst (30–50 μM [Fe(TPP)]total).
The experiments have some uncertainty in the catalyst speciation with time due to the
overlapping absorbance features, low concentrations of the different Fe(TPP) species, and the
uncertainties in the experimental ε values (in part due to high air sensitivity at low concentrations
and instability of the materials). Therefore, we sought a model that fit the general trends in
catalyst speciation rather than the exact concentration profiles (Figure 6). The fits are good for
the [Fc*+] time courses and agree with the general trends of catalyst speciation. Notably, the
•‒
model fits well to changes in initial [FeII(TPP)] and [FeIII(TPP)(O2 )] concentrations across a variety
of temperatures and is representative of the temperature-dependent O2 binding equilibrium.
Across the whole temperature range, the model also correctly predicts that [FeIII(TPP)]OTf is the
predominant catalyst species at the end of the reaction, even when excess Fc* is used.
The accuracy of the model is evidenced by the remarkable agreement with experimental
data for both the thermodynamic and kinetic parameters (Table 1). The modeled values for ∆H°ET,
∆H°O2 and ∆S°ET very closely match the directly determined experimental values from Sections I
and II above. The modeled value for ∆S°O2 is somewhat less negative than the experimental value,
which is similar to typical entropies for O2 binding in polar organic solvents.48 The inconsistencies
between modeled and experimental ∆S°O2 could be due to an unrecognized catalyst complex in
the spectral fitting. Although we cannot rule out such a contribution, we believe the higher than
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expected modeled ∆S°O2 may instead be related to the lower than expected ∆S‡ , which would
PT
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7
usually be a positive value for a bimolecular reaction. Thus, the ∆G‡ of 11 kcal mol-1 at 298 K
PT
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derived from the modeled may be an underestimation. At room temperature, the model predicts
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ln(k ) to be between 7.1-7.5 and ln( ) to be between 5.3-7.5. Under electrochemical
퐾O
PT
2
conditions, where k is the product of
and kPT, the model predicts kcat to be between 2-20 
퐾O
cat
2
105 M–2 s–1, which is within error of the experimental kcat value of (3.2‒6.4)  105 M-2 s-1. This
agreement in rate constants at 298 K could also be suggestive of a balancing of errors in the
modeled ∆S°O2 and ∆S‡ . Although ∆S°O2 and ∆S‡PT are not significantly correlated according to
PT
the COPASI statistical analysis, there is evidence that these parameters are not all known
completely independently from one another. We can achieve similarly strong fits of the data with
different parameter values when they are tightly constrained the to the experimental results (see
SI, Section 6).
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Figure 6. Data (points) and fits (lines) for three stopped-flow ORR time courses (0.3 μM Fe(TPP),
50 mM pTsOH, 0.33 mM O2, 3 mM Fc*, varied temperatures: a) 253 K, b) 283 K, c) 303 K. For each
temperature column, the top figure shows the Fc*+ concentrations with an expansion for the
Fe(TPP) speciation below.
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•‒
VI. Computational Analysis of the Barrier for protonation of FeIII(TPP)(O2 ) by pTsOH:
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that computational analysis to evaluate the thermochemistry and kinetics of protonation of
•‒
FeIII(TPP)(O2 ) by pTsOH, the acid used in this study.
•‒
The thermochemistry and kinetics of FeIII(TPP)(O2 ) protonation are dictated by
interactions between the proton source, solvent, and the superoxide adduct.32 As such, DFT
analysis was performed to quantify the energetics of interactions between explicit pTsOH, pTsO–,
and DMF molecules in DMF solvent. Calculations show that the free energy of the pTsOH•••DMF
pair is stabilized by 1.6 kcal mol-1 relative to its constituent parts (separate pTsOH and DMF
molecules). Formation of the [pTsOH•••pTsO]– homoconjugate is even more stable, favored by
5.9 kcal mol-1 relative to separate pTsOH and pTsO– molecules; however, the experimental
conditions ([DMF] >>> [pTsO–]) favor the pTsOH•••DMF heteroconjugate. Therefore, the
•‒
thermochemistry and kinetics of proton transfer to FeIII(TPP)(O2 ) were investigated using
pTsOH•••DMF as the proton donor.
•‒
The computed free energy profiles for protonation of FeIII(TPP)(O2 ) (1) by pTsOH•••DMF
are shown in Figure 7a, wherein all energetics are referenced to 1. The free energy profile for
protonation by pTsOH•••DMF first features the formation of a pre-association complex, 2, in
•‒
which the acid approaches FeIII(TPP)(O2 ) (1). Formation of 2 involves solvent reorganization and
the establishment of a three-center site and has an overall energetic penalty of 8.5 kcal mol-1.
This energetic cost results from unfavorable entropic (–TS = 10.7 kcal mol-1 at 298 K) and
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solvation terms (7.8 kcal mol-1) that are only partially balanced by being electronically favorable
(–9.9 kcal mol-1). Proton transfer then yields the perhydroxyl (3) species, which was calculated to
be 9.8 kcal mol-1 uphill from 1. This result indicates that the proton transfer is unfavorable and is
consistent with the experimental data.
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Figure 7. a) Computed free energy profile for association (Kassc) and protonation of
⦁–
FeIII(TPP)(O2 ) by the pTsOH•••DMF heteroconjugate pair. b) Optimized structure for the proton
transfer transition state (kPT) and the corresponding O-H bond lengths for the proton transfer
•‒
donor (pTsOH), solvent (DMF), and acceptor FeIII(TPP)(O2 ). All distances in Ångstroms.
To further understand the kinetics of proton transfer, the barrier for protonation from
the heteroconjugated pair pTsOH•••DMF was sought. As shown in Figure 7, the transition state
(TS*) involves proton transfer from the pre-association complex (2) and has an overall barrier of
•‒
22.0 kcal mol-1 relative to FeIII(TPP)(O2 ). This calculated barrier is an upper bound, as it does not
compensate for the errors in computing sulfonate group energetics. Taking into account these
errors as discussed in the SI, the estimated lower bound of this barrier is about 17.0 kcal mol-1.
The energetic contributions for the energy change between 2 and TS* show that the transition
state energy is enthalpically driven. Additionally, at the transition state, the proton is closer to
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•‒
the acceptor FeIII(TPP)(O2 ) than the donor pTsOH•••DMF, as evidenced by the O-H bond
lengths in the optimized structure (Figure 7b). The kinetic isotope effect for this transition state
was calculated using pTsOD•••DMF as the proton donor, affording a KIE of 0.94 (see section 8
SI).
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Discussion
I. The mechanism of ORR catalysis by Fe(TPP)
The results of the spectroscopic, electrochemical, and computational experiments
described above support the mechanism illustrated in Scheme 1. In this mechanism, Fc* or an
electrode rapidly reduce [FeIII(TPP)]OTf to the ferrous porphyrin, FeII(TPP). FeII(TPP) then binds
•‒
O reversibly ( ) to form the ferric superoxide, FeIII(TPP)(O2 ). The rate-determining step is
퐾O
2
2
protonation of FeIII(TPP)(O2 ) (kPT) to form the perhydroxyl-iron(III) complex, [FeIII(TPP)(O2H⦁)]+.
•‒
The perhydroxyl complex is rapidly reduced and protonated under the catalytic conditions to
produce two equivalents of H2O and restart the catalytic cycle. The evidence in support of this
mechanism is discussed below.
Reduction of [FeIII(TPP)]OTf to FeII(TPP), the first step in the proposed catalytic cycle, was
found to be a fast pre-equilibrium between [FeIII(TPP)]OTf and the reductant in both
electrochemical and spectroscopic measurements. The cyclic voltammograms during
electrocatalytic O2 reduction fit well to a mechanism in which rapid pre-equilibrium electron
transfer from the electrode ([FeIII(TPP)]+ + e– FeII(TPP)) is followed by a rate limiting catalytic
⇌
step, kobs (an EC mechanism). With Fc* as a chemical reductant, this pre-equilibrium is
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thermodynamically uphill at standard state (∆G°ET = +1.01(8) kcal mol-1). Spectroscopic studies of
the reduction of [FeIII(TPP)]OTf by Fc* at different temperatures have yielded the ∆H°ET and ∆S°ET
for this equilibrium (Results Section I). Those values are consistent with the electrochemically
measured difference in reduction potentials between [Fe(TPP)]+/0 and Fc*+/0 at ambient
temperatures. In the room temperature stopped-flow kinetics, the initial spectra show almost
complete reduction of [FeIII(TPP)]OTf only because the starting concentration of Fc*+ is very small
~0 mM and the [Fc*]/Fc*+ ratio is very large. As catalysis progresses, however, the ratio of Fc* to
Fc*+ decreases, quickly shifting this equilibrium so that [FeIII(TPP)]+ is the resting state of the
catalyst for much, if not most, of the reaction (Figures 6 and S34-S49).
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Spectroscopic experiments have also quantified the rapid pre-equilibrium O2 binding to
FeII(TPP), which is proposed to follow initial electron transfer. The distinct optical spectrum of the
•‒
superoxide complex FeIII(TPP)(O2 ) enabled measurements of the equilibrium constants for
binding ( ) at low temperatures in the absence of reductant and acid. Low temperatures are
퐾O
2
•‒
required to avoid the decomposition of FeIII(TPP)(O2 ) to the well-known -oxo dimer.42 The
measured ∆H° and ∆S° could then be used to predict values of
at the higher temperatures
퐾
O2
O2
O2
of catalysis. The -oxo dimer is not formed under the catalytic conditions because of the presence
of strong acid, as indicated reaction of independently prepared -oxo dimer with pTsOH (see SI,
Section 5.4) and by prior reports.49,50 Consistent with these equilibrium measurements,
•‒
FeIII(TPP)(O2 ) is observed as the predominant initial catalyst resting state in catalytic stopped-
flow experiments at low temperatures (e.g. 253 K). Under such conditions, [FeIII(TPP)]OTf is
•‒
completely converted to FeIII(TPP)(O2 ) within 0.1 s (Figure S29). As the reaction proceeds, both
[O2] and the ratio of Fc*:Fc*+ decrease, leading to a change in resting state from reduced
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porphyrin intermediates to [FeIII(TPP)]+ (Figures 6 and S34-S49).
5
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8
9
The kinetics of the electrocatalysis are not complicated by the shifts in the pre-equilibrium
reduction or O2-binding steps observed in the stopped flow data. The catalytic voltammograms
are fit using FOWA, which yields rate constants for reactions catalyzed by the reduced FeII(TPP)
catalyst resting state. Additionally, because the current is analyzed at the ‘foot’ of the wave, the
concentration of substrates in the reaction-diffusion layer is the same as the bulk solution (true
pseudo-first order conditions). For these reasons, the electrocatalytic kinetics yield a simple third
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order rate law, as previously reported,31 obtained by removing the
, [Fc*], and [Fc*+] terms
퐾ET
from the complex rate law required for the stopped-flow data (eq 9). Despite the complex rate
law using Fc* as a chemical reductant, the similarities between the kcat values obtained
electrochemically and chemically confirm that electron transfer kinetics are not involved in the
rate limiting steps and remain a fast pre-equilibrium step under all conditions. Furthermore, the
first order dependence on O2, pTsOH and [Fe(TPP)] exclude bimolecular pathways that involve
peroxo (Fe-O2-Fe) or -oxo (Fe-O-Fe) dimeric intermediates, pathways that are observed in the
absence of strong acids.51–54
Catalytic turnover requires the addition of four protons and three more electrons to
•‒
FeIII(TPP)(O2 ), ultimately forming water and [FeIII(TPP)]OTf. A number of initial steps could be
imagined for this transformation, but only those which have an acid dependence were
considered since the rate of the ORR is first order in [pTsOH]. Within that constraint, the two
possible pathways for the reaction of the superoxide complex involve i.) proton transfer (PT) from
pTsOH to form perhydroxyl or ii.) concerted proton-coupled electron transfer (CPET) with pTsOH
and a reductant to give a hydroperoxo complex. While option ii. should be thermodynamically
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advantageous, it is formally a termolecular CPET reaction and thus is kinetically challenging in the
absence of a favorable pre-association with the proton donor.55,56 As DFT calculations show that
7
•‒
pre-association of pTsOH with FeIII(TPP)(O2 ) is uphill by 8.5 kcal mol-1 (Figure 7), pathway ii. is
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•‒
unlikely, and rate-limiting proton transfer from pTsOH to FeIII(TPP)(O2 ) is the more plausible
mechanism. Previous studies of ORR by Fe(TPP) under different conditions – using HClO4 as the
acid and in the presence of excess chloride – found a zero order dependence of turnover
frequency on the concentration of Fc*.31 In the model used successfully here (Scheme 2), the
concentration of the electron donor appears only in the first pre-equilibrium. The rate
determining step cannot be CPET because that would require a molecule of reductant.
The strongest support for the proposed mechanism is the close agreement between the
equilibrium parameters determined by COPASI-fits to the stopped flow data and the directly
measured equilibrium values of ∆H°ET/∆S°ET and ∆H°O2/∆S°O2. In addition, the idealized third-
order rate constant predicted by the COPASI-fit, (2-20)  105 M-2 s-1, is in excellent agreement
with the value determined by electrocatalysis, kcat
=
(3.2
–
6.4)

105
M-2 s-1.
The computational conclusions qualitatively agree with experimental results and show
that proton transfer is the turnover-limiting step. The barrier extracted from the COPASI fits,
∆G‡PT  11 kcal mol-1, is significant but smaller than the computed barrier. As noted above, the
modeled ∆G‡ may be low due to an underestimation of the ∆S‡ , which would bring the
PT
PT
experimental and computed values into closer agreement. We note that the complexity of the
COPASI kinetic modeling and the challenges of computing sulfonate groups make close
agreement between experiment and theory unlikely. Still, good agreement between theory and
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