Facile and Reversible Formation of Iron(III)–Oxo–Cerium(IV) Adducts from Nonheme Oxoiron(IV) Complexes and Cerium(III)

Article abstract of DOI:10.1002/anie.201704322

Ceric ammonium nitrate (CAN) or Ce^IV(NH₄)₂(NO₃)₆ is often used in artificial water oxidation and generally considered to be an outer-sphere oxidant. Herein we report the spectroscopic and crystallographic characterization of [(N4Py)Fe^III-O-Ce^IV(OH₂)(NO₃)₄]⁺ (3), a complex obtained from the reaction of [(N4Py)Fe^II(NCMe)]²⁺ with 2 equiv CAN or [(N4Py)Fe^IV=O]²⁺ (2) with Ce^III(NO₃)₃ in MeCN. Surprisingly, the formation of 3 is reversible, the position of the equilibrium being dependent on the MeCN/water ratio of the solvent. These results suggest that the Fe^IV and Ce^IV centers have comparable reduction potentials. Moreover, the equilibrium entails a change in iron spin state, from S=1 Fe^IV in 2 to S=5/2 in 3, which is found to be facile despite the formal spin-forbidden nature of this process. This observation suggests that Fe^IV=O complexes may avail of reaction pathways involving multiple spin states having little or no barrier.

Full text of DOI:10.1002/anie.201704322

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Accepted Article
Title: Facile and reversible formation of Fe(III)-O-Ce(IV) adducts from
nonheme oxoiron(IV) complexes and Ce(III)
Authors: Apparao Draksharapu, Waqas Rasheed, Johannes E. M. N.
Klein, and Lawrence Que
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To be cited as: Angew. Chem. Int. Ed. 10.1002/anie.201704322
Angew. Chem. 10.1002/ange.201704322
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10.1002/anie.201704322
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Facile and reversible formation of Fe(III)–O–Ce(IV) adducts from
nonheme oxoiron(IV) complexes and Ce(III)
Apparao Draksharapu, Waqas Rasheed, Johannes E. M. N. Klein, Lawrence Que, Jr.*
Dedicated to Eckard Münck in celebration of his outstanding career in bioinorganic chemistry
Abstract: CAN or CeIV(NH4)2(NO3)6 is often used in artificial water
oxidation and generally considered to be an outer-sphere oxidant.
Herein we report the spectroscopic and crystallographic
characterization of [(N4Py)FeIII–O–CeIV(OH2)(NO3)4]+ (3), a complex
obtained from the reaction of [(N4Py)FeII(NCMe)]2+ with 2 equiv. CAN
or [(N4Py)FeIV=O]2+ (2) with CeIII(NO3)3 in MeCN. Surprisingly, the
formation of 3 is reversible, the position of the equilibrium being
dependent on the MeCN/water ratio of the solvent. These results
suggest that the FeIV and CeIV centers have comparable reduction
potentials. Moreover, the equilibrium entails a change in iron spin
state, from S = 1 FeIV in 2 to S = 5/2 in 3, which is found to be facile
despite the formal spin forbidden nature of this process. This
observation suggests that FeIV=O complexes may avail of reaction
pathways involving multiple spin states having little or no barrier.
[(TMC)FeIV(O)(NCMe)]2+ (1) with ScIII(OTf)3.[8] The corresponding
reaction with CrII(OTf)2 affords the FeIII–O–CrIII adduct.[9] Here we
report the reaction of [(N4Py)FeIV(O)]2+ (2, N4Py = N,N-bis(2-
pyridylmethyl)-N-bis(2-pyridyl)methylamine) with CeIII(NO3)3 to
form [(N4Py)FeIII–O–CeIV(OH2)(NO3)4]+ (3) (Scheme 1). Unlike
observed for the Sc- and Cr-adducts of 1, formation of the Ce-
adduct of 2 can be reversed, simply by adjusting the MeCN/water
ratio of the solvent. The existence of such an equilibrium
demonstrates facile electron exchange between the Fe and Ce
centers, despite a formally forbidden change of spin state from S
= 1 FeIV to S = 5/2 FeIII. Parallel changes are also observed for
[(BnTPEN)FeIV(O)]2+ (4) (Scheme 1).
Scheme 1. Equilibria between (L)FeIV(O) complexes, 2 (L = N4Py) and 4 (L =
Nature uses a Mn4CaO5 cluster to perform efficient water
oxidation,[1] a fact that has inspired the development of artificial
photosystems based on ruthenium, iridium, manganese, iron,
cobalt and copper.[2] In these systems, high-valent metal-oxo
species are often proposed to be the active species or precursors
thereof, which are commonly generated by ceric ammonium
nitrate (CAN) as a sacrificial outer-sphere oxidant.[3] However, a
fleeting species has recently been identified by Lloret and Costas
in a catalytic water oxidation system consisting of [(mcp)FeIV(O)]2+
BnTPEN), and (L)FeIII–O–CeIV(NO3)4 species, 3 (L = N4Py) and 5 (L = BnTPEN).
Treatment of [(N4Py)FeII(NCMe)](OTf)2 in MeCN with 2
equiv. CAN in 10 µL water generates a brown species 3 with a
shoulder near 500 nm (Figures 1a and S1), a result clearly distinct
from that reported by Nam showing the formation of
[(N4Py)FeIV=O]2+ (2, λmax 696 nm) in quantitative yield upon
addition of 4 equiv. CAN to 1 in 1:3 (v/v) H2O/MeCN.[11] Complex
3 is also formed upon addition of 4.5 equiv. Ce(NO3)3 to 2 in
MeCN, resulting in the immediate loss of absorbance at 696 nm
and concomitant increase at ca. 500 nm to form 3 with an
isosbestic point observed at 570 nm (Figure 1b).
(mcp
=
N,N'-dimethyl-N,N'-bis(2-pyridylmethyl)-1,2-cis-
diaminocyclohexane) and CAN in water at pH 1.[4] Based on cryo-
spray mass spectrometry and resonance Raman spectroscopy,
this intermediate has been formulated as [(mcp)(O)FeIV–O–
CeIV(NO3)3]+, where CeIV acts as both a Lewis acid and an inner-
sphere oxidant.
The redox-inactive Ca2+ ion in the oxygen evolving complex
(OEC) of photosystem II is essential for O–O bond formation and
ultimately dioxygen generation.[5] One proposed role for the Ca2+
is to increase the potential of the manganese oxido cluster.[6] This
notion is also in agreement with the observation that Lewis acids
such as ScIII(OTf)3 increases the FeIV/III redox potential of
nonheme FeIV=O complexes.[7] That ScIII interacts directly with the
FeIV=O unit is supported by the crystal structure of [(TMC)FeIII–
O–ScIII(OTf)4(solvent)] (TMC
=
1,4,8,11-tetramethyl-1,4,8,11-
tetraazacyclotetradecane)
formed in the reaction of
[a]
Dr. Apparao Draksharapu, Waqas Rasheed, Dr. Johannes E. M. N.
Klein, Prof. Dr. Lawrence Que, Jr.
Figure 1. UV-Vis spectral changes in (a) the reaction of [(N4Py)FeII(NCMe)]2+
(2.5 mM in MeCN, dashed line) with 2 equiv. CAN in 10 µL H2O (solid line just
above dashed line) followed by titration with water (final concentrations 2.8 M,;
3.9 M,; 5 M,; 8.3 M), resulting in increasing absorbance at 696 nm (# =
vibrational overtone of water) and (b) the titration of 2 (A696 nm = 1; 2.5 mM in
MeCN) with CeIII(NO3)3 in MeCN (0.25, 0.5, 0.75, 1, 1.5, 2, 3, 4.5 equiv.),
resulting in diminishing absorbance at 696 nm.
Department of Chemistry and Center for Metals in Biocatalysis,
University of Minnesota, Minneapolis, Minnesota 55455, United
States
E-mail: larryque@umn.edu
Supporting information for this article is given via a link at the end of
the document.
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ESI-MS analysis of 3 in MeCN reveals two major peaks at
m/z 827.1 and 914.0 with masses and isotope patterns
corresponding to [(N4Py)FeIII–O–CeIV(NO3)3(X)]+ ions with X =
NO3 and OTf (Figures S5-S7). In support of the iron(III) oxidation
state assignment, 3 exhibits at 4 K broad EPR signals with g = 8.9
and 4.3, which are characteristic of a rhombic S = 5/2 FeIII species
(Figure S9). The resonance Raman spectrum of 3 (exc 514.5 nm)
shows an enhanced band at 707 cm-1 (Figure 2a) that shifts to
676 cm-1 upon 18O labeling, the 31-cm-1 downshift being in
excellent agreement with the calculated value based on Hooke’s
law of 31 cm-1 for a diatomic Fe–O oscillator. As the Ce atom is
much heavier than the iron atom, the former would not be
expected to contribute to the Fe–O stretch. For comparison, this
vibration has a slightly higher frequency than that found by
Borovik for [(H3buea)FeIII(O)]2- (671 cm-1, H3buea = tris[(N'-tert-
butylureaylato)-N-ethylene]aminato)trianion).[12]
both of which have high-spin ferric centers (Table S5). Lastly, 3
has an Fe···Ce distance of 3.825(5) Å, which is longer than those
found for other FeIII–O–M complexes, e.g.  3.61 Å for M =
FeIII,[9b],[14] and 3.64 for M = ScIII,[8] reflecting the larger ionic radius
of the CeIV center.
The cerium atom in 3 is connected to the iron half of the molecule
via the oxo bridge. The Ce–-O bond length is 2.013(2) Å, much
shorter than the 2.418(2)-Å distance observed for the Ce–OH2
bond in 3, which is within the range of reported aquacerium(IV)
complexes.[15] However the Ce–-O bond of 3 is longer than the
CeIV=O bonds observed for (H2O)(LOEt)2CeIV=O (1.857(3) Å)
(LOEt
=
CpCo{P(O)–(OEt)2}3)
and
[(N(CH2CH2NSitBuMe2)3)CeIV=O]– (1.902(2) Å).[16] In addition, the
Ce center has four 2-bound nitrate ligands with an average Ce–
O bond length of 2.52 Å, which is consistent with the 2.508(7)-Å
distance observed in CAN.[17]
Figure 2. (a) Resonance Raman spectra (exc 514.5 nm) of 3 from the reaction
of 5 mM [(N4Py)FeII(NCMe)]2+ in MeCN with 2.2 equiv. CAN in 10 µL H2O
(black) and the corresponding 18O-labeled complex generated from 2.2 equiv.
CAN in 10 µL H218O (red). # represents a MeCN deformation. (b) Unfiltered Fe
K-edge EXAFS spectrum (dotted line) and best fit (solid line) of 3. Inset:
corresponding unfiltered k-space data (dots) and best fit (line).
Figure 3. ORTEP plot for [(N4Py)FeIII–O–CeIV(OH2)(NO3)4]+ (3) with 50%
probability ellipsoids (H atoms omitted for clarity). Selected distances
(Å): Fe(1)−O(1),
1.825(2);
Fe(1)−N(1),
2.162(3); Fe(1)−N(2),
2.084(3); Fe(1)−N(3), 2.072(3); Fe(1)−N(4), 2.108(3); Fe(1)−N(5), 2.117(3);
Ce(1)−O(1), 2.013(2); Ce(1)−O(14), 2.418(3); Ce(1)−O(nitrate), 2.518(3);
Ce(1)•••Fe(1), 3.8247(5). Fe(1)−O(1)−Ce(1) = 170.3(1)º. CCDC 1546446.
To shed further light into its nature, X-ray absorption
spectroscopy (XAS) studies have been carried out on frozen
solutions of 3. The Fe K-edge for 3 is observed at 7122.7 eV
(Figure S12), a value comparable to those of [(TMC)FeIII–O–
ScIII(OTf)4(L)] (7122.6 eV)[8b] and other ferric complexes.[13]
Complex 3 exhibits a single pre-edge feature (assigned to a 1s 
3d transition) at 7112.9 eV with an area of 13.7 units, consistent
with values found for (-oxo)diferric complexes.13c The Fourier
transformed EXAFS spectrum of 3 shows prominent features at
R + Δ ~ 1.7 and 3.4 Å (Figure 2b) corresponding to 1 N/O scatterer
at 1.84 Å (the oxo bridge), 5 N/O scatterers at 2.13 Å (N4Py N-
atoms) and a Ce scatterer at 3.8 Å indicating the presence of a
nearly linear FeIII–O–CeIV core (see Table S5 for further details).
The thermal stability of 3 led to the growth of diffraction
quality crystals by slow diffusion of ether into a solution of 2
(derived from the oxidation of [(N4Py)FeII(NCMe)](ClO4)2) in the
presence of 5 equiv. CeIII(NO3)3 in acetonitrile at -20 ºC. The
crystallographic analysis (Figure 3, Table S5) confirms the results
from XAS analysis. There is a nearly linear FeIII–O–CeIV unit with
an Fe–O–Ce angle of 170.3º, similar to the Fe–O–Fe angle (172º)
found in [(N4Py)FeIII–O–FeIII(N4Py)]4+.[14] Its Fe–O bond length of
1.825 Å is significantly longer than the 1.64-Å Fe=O bond of 2 but
comparable to those of [(N4Py)2FeIII2(O)]4+ (1.8034(10) Å) and
[(N4Py)FeIII(OCH3)]2+ (1.77 Å).[14] In addition, the average Fe–NPy
bond length of 2.11 Å found for 3 is comparable to those of
[(N4Py)FeIII(OCH3)]2+ (2.11 Å) and [(N4Py)2FeIII2O]4+ (2.13 Å),
Unexpectedly, subsequent experiments revealed 2 and 3 to
be involved in a reversible equilibrium. Figure 1a shows the effect
of titrating a MeCN solution of 3 with water, where the 696-nm
band associated with 2 was fully formed at a water concentration
of 7.5 M, with an isosbestic point at 580 nm. Subsequent dilution
of this sample with MeCN shifted the equilibrium back in favor of
3. This conclusion was corroborated by ESI-MS experiments that
showed the disappearance and re-appearance of the ions
associated with
3 under the different solvent conditions
concomitant with the appearance and disappearance of 2 (Figure
S6). Analogous experiments with [(BnTPEN)FeIV=O]2+ (4)
(BnTPEN
=
N-benzyl-N,N',N'-tris(2-pyridylmethyl)-1,2-
diaminoethane) and CeIII(NO3)3 afforded [(BnTPEN)FeIII–O–
CeIV(OH2)(NO3)4]+ (5), with comparable properties as 3 (Figures
S3, S4, S6, S9-S11, S13, S15). In contrast to 2 and 3, the
conversion of 4 to 5 required only 2 equiv. CeIII(NO3)3, but more
water (10 M) was needed to shift the equilibrium between 4 and 5
back in favor of 4. The equilibria demonstrated between 2 and 3
and between 4 and 5 represent the first examples of reversible
inner-sphere electron transfer between CeIII and an FeIV=O
complex.
The existence of such equilibria suggests that the FeIV/III
potentials of 2 and 4 and the CeIV/III potentials of 3 and 5 are not
very different from each other. Where each equilibrium lies under
the different conditions depends on the respective ligand
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environments of the two metal centers involved. Based on our
accumulated observations, 4 would appear to have a somewhat
higher redox potential than 2, as less Ce(NO3)3 is required to form
5 from 4 but more water is needed to convert 5 back to 4. That 4
is more oxidizing than 2 is consistent with conclusions derived
from comparative electrochemical data (Table 1). On the other
hand, the observed shifts in the above equilibria from dry MeCN
to wet MeCN reflect an increase in the CeIV/III potential (Figure
S20) as the nitrates bound to CeIV in 3 and 5 are progressively
replaced by water.
To put the reactions of 2 and 4 with CeIII(NO3)3 into a
broader context, we added CeIII(NO3)3 to [(TMC)FeIV=O]2+ (1) in
acetonitrile but observed no reaction even with a 20-fold excess
of CeIII or with added nitrate. These observations suggest that 1
is a significantly weaker oxidant than 2 or 4, consistent with the
observed reactivities of these three complexes with respect to H-
atom and O-atom transfer.[18],[19] On the other hand, Bakac has
reported that [(H2O)5FeIV=O]2+ irreversibly oxidizes CeIII(NO3)3
(CeIV/III 1.45 V vs. SCE) in 1 M HClO4.[21] Based on this
comparison, the relative oxidizing power of these FeIV=O
complexes decreases in the order: [(H2O)5FeIV=O]2+ >> 4 > 2 > 1.
Fukuzumi and Nam have studied the fundamental electron
transfer properties of ferrocene derivatives to nonheme FeIV=O
complexes and applied the Nernst equation under equilibrium
conditions to determine the reduction potentials of 1, 2, and 4 to
be about 0.4-0.5 V vs SCE (Table 1).[20] While this approach
allowed the redox potentials of these three nonheme FeIV=O
complexes to be compared for the first time, a potential concern
is whether a true equilibrium can be established under the
experimental conditions, because of the instability of the reduced
FeIII–O– species.[22] In contrast, the reduction of 2 and 4 by
CeIII(NO3)3 reported here is facile and reversible, suggesting that
2 and 4 must be relatively close to CeIII(NO3)3 in redox potential
in MeCN solvent. As shown in Figure S21, the CeIV/III potential of
CeIII(NO3)3 can be tuned between 1-1.3 V vs. SCE by adjusting
the Ce:NO3 ratio.[23] At a 1:4 Ce:NO3 ratio, 1 equiv. CeIII is
sufficient to reduce 2 and 4 to 3 and 5, respectively, so 2 and 4
must have redox potentials higher than the CeIII species present
in the solution (E1/2 > 1 V vs SCE), in agreement with values
obtained by spectropotentiometric oxidation of corresponding
(N4Py)FeIII–OH and (BnTPEN)FeIII–OH complexes (Table 1).[18]
Because Fukuzumi and Nam have shown that a Lewis acid
such as ScIII(OTf)3 can interact with the FeIV=O unit to increase
the FeIV/III redox potential by 0.84 V,[7] the higher potentials
deduced for 2 and 4 from the CeIII(NO3)3 titration experiments may
reflect the Lewis acid effect of Ce(III). We have thus investigated
the effect of 1,1'-diacetylferrocene (Ac2Fc+/0), another reductant
that has a similarly high redox potential (0.89 V vs. SCE) but
cannot act as a Lewis acid. As shown in Figure S22, 1 equiv. 1,1'-
diacetylferrocene can in fact reduce 2 and 4, corroborating the
high potentials deduced for 2 and 4 and excluding the Lewis acid
argument to rationalize the high potentials deduced from the
CeIII(NO3)3 titrations. However, despite having a lower redox
potential than CeIII(NO3)3, Ac2Fc reduces 2 and 4 much more
slowly, requiring 20 min to reduce 4 completely at 20 °C and ten
times longer for 2. This comparison suggests that the Lewis acid
does not affect the thermodynamics of the reduction but instead
enhances its kinetics presumably by proceeding via a mechanism
different from that for Ac2Fc reduction. The latter very likely occurs
by outer-sphere electron transfer, but the Ce(III) reductant offers
the possibility of an inner-sphere pathway for electron transfer by
coordinating to the FeIV(O) unit, an intriguing notion we will further
investigate in future work.
Table 1. Redox properties of select nonheme FeIV(O) complexes.[a]
1
2
4
Refs
Reaction with CeIII(NO3)3 in
MeCN
No rxn
yes
yes
This
work
[18], [19]
Ep,c by cyclic voltammetry
in dry MeCN [b]
+0.08
---
-0.13
-0.03
[18]
E1/2 by spectropotentiometry
in wet MeCN [c]
+1.30
+1.47
[20]
[7]
Ered from ferrocene titrations [d]
+0.39
---
+0.51
+1.35
+0.49
---
[FeII(bpy)3]2+ titrations
in the presence of ScIII [e]
[a] All potential values are vs. SCE. [b] Ep,c values obtained from cyclic
voltammetry of FeIV=O complexes in anhydrous MeCN. [c] From
spectropotentiometry starting with Fe(II) complexes in MeCN with water. [d]
From titrations of FeIV=O complexes with ferrocenes. [e] From titration of 2 with
[FeII(bpy)3]2+ in the presence of >1000 equiv. ScIII(OTf)3 in MeCN.
The effect of added nitrate ion has also been investigated,
as the crystal structure of 3 shows the CeIV center to have four 2-
bound nitrate ligands (Figure 3) and the CeIII(NO3)3 reductant
introduced only has three nitrates. Indeed, with the addition of one
equiv. Bu4N(NO3), both 2 and 4 can be fully converted to 3 and 5,
respectively, with just one equiv. CeIII(NO3)3 (Figures 4 and S16),
demonstrating that the coordination of four nitrates to the CeIV
center stabilizes the FeIII–O–CeIV species. We interpret these
results to indicate a decrease in the CeIV/III potential upon binding
of the fourth nitrate, a notion confirmed by cyclic voltammetry
(Figure S21). Importantly, the requirement for only one equiv.
CeIII(NO3)3 shows that the speciation in the solution corresponds
precisely to what has been demonstrated by X-ray
crystallography (Figure 3).
In this work, we have demonstrated the synthesis of FeIII–
O–CeIV complexes from the reactions of their FeIV=O precursors
with CeIII(NO3)3 in what turns out to be a facile and reversible
transformation. The position of this equilibrium depends on the
number of nitrate ligands on the Ce center as controlled by the
MeCN/H2O ratio in the solvent, which tune the CeIV/III potential.
Figure 4. (a) UV-Vis spectral changes observed upon titration of 1 mM 2 (A696-
= 1) and 1 equiv. Bu4N(NO3)in MeCN with 0.1-equiv. aliquots of CeIII(NO3)3
nm
in MeCN. (b) Plot of A(696 nm) as a function of CeIII(NO3)3 concentration.
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These observations imply that the FeIV=O centers of 2 and 4 have
reduction potentials higher than that of the CeIII reductant in
MeCN (~1 V vs SCE). More importantly, the existence of 3 and 5
supports the formation of analogous Fe–O–Ce adducts in Fe-
catalyzed water oxidation as proposed by Lloret-Fillol and
Costas.[4] Our work corroborates the assignment of an observed
Raman band at 677 cm-1 in Fe-catalyzed water oxidation to an
Fe–O–Ce intermediate that facilitates formation of the O–O bond,
suggesting that CAN may also play such a role in water oxidation
reactions catalyzed by other metal complexes. Lastly, despite
being a formally spin-forbidden transformation, the conversion of
an S = 1 FeIV=O complex to the corresponding (S = 5/2 FeIII)–O–
CeIV species upon addition of CeIII is remarkably facile, being
instantaneous even at -40 °C (Figure S23). This observation
suggests that the barrier for such spin crossover events at FeIV=O
centers is quite low, an aspect discussed in a recent review by
Harvey on computational insights into spin forbidden reactions.[24]
Furthermore, this Ce-based conversion bears some analogy to
the conversion of an S = 1 FeIV=O complex to an S = 5/2 FeIII–OH
species in an H-atom transfer (HAT) reaction, with CeIII and H•
performing similar functions and is fundamentally different from
metal-coupled electron transfer (MCET) reactions of FeIV=O
centers in the presence of Lewis acids where the electron does
not originate from the Lewis acid, but rather an external reducing
equivalent.[25] Assuming this comparison to be valid, little or no
barrier may be expected for the S = 1 FeIV=O reactant to undergo
spin-crossover as it initiates HAT. We therefore conjecture that,
for S = 1 FeIV=O centers, it is not the ability to undergo spin
crossover that dictates which spin surface is relevant to its HAT
reactivity, but rather the gap between the S = 1 ground state and
the S = 2 excited state, as invoked in the Two-State Reactivity
model of Shaik[10] and recently demonstrated by the correlation of
HAT rates with spin state splitting energies for a series of eleven
FeIV=O complexes supported by the tetramethylcylam ligand.[26]
Keywords: water oxidation • iron oxo • CAN • inner-sphere
electron transfer • Fe–O–Ce
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Experimental Section
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This work was supported by a grant from the U. S. National
Institutes of Health (GM-38767 to L.Q.). XAS data were collected
at the Stanford Synchrotron Radiation Lightsource, which is
supported by the U.S. DOE under Contract No. DE-AC02-
76SF00515. Use of Beamline 7−3 is supported by the DOE Office
of Biological and Environmental Research and the National
Institutes of Health, National Institute of General Medical
Sciences (including P41GM103393). J.E.M.N.K. thanks the
Alexander von Humboldt Foundation for a Feodor Lynen
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10.1002/anie.201704322
Angewandte Chemie International Edition
COMMUNICATION
Entry for the Table of Contents
COMMUNICATION
Cerium caught in the redox act!
Apparao Draksharapu, Waqas
Rasheed, Johannes E. M. N. Klein,
Lawrence Que, Jr.*
CeIII reacts with FeIV(O) complexes in
MeCN to form inner-sphere FeIII–O–
CeIV adducts, a transformation easily
reversed by adding water. This facile
equilibrium is tuned by the nature of
the ligands on the Ce center, and its
existence shows that the CeIV/III and
FeIV/III potentials must fall within a
range of 0.9 - 1.1 V vs SCE.
Page No. – Page No.
Facile and reversible formation of
FeIII–O–CeIV adducts from nonheme
oxoiron(IV) complexes and Ce(III).
This article is protected by copyright. All rights reserved.