Intermolecular versus intramolecular electron-/atom- (Cl{radical dot}) transfer in heme-iron and copper pyridylalkylamine complexes

Article abstract of DOI:10.1016/j.ica.2007.08.016

While outer-sphere electron-transfer reactions come with a firm experimental and theoretical basis, less is known about the redox reactions occurring by atom-transfer. In the present study, relevant reactions occur upon mixing, (F₈)Fe^III(Cl) (F₈ is tetrakis(2,6-difluorophenyl)porphyrinate) with a series of [(L)Cu^I]⁺ complexes (L are tripodal tetradentate pyridylalkylamine ligands varying in effective chelate ring sizes: TMPA, PMEA, PMAP) to form (F₈)Fe^II and [(L)Cu^II(Cl)]⁺. The electron-/atom- (Cl^{{radical dot}}) transfer process is characterized by electrochemical measurements as well as UV-Vis, ¹H NMR, and EPR spectroscopies. Stopped-flow UV-Vis spectroscopy in THF indicate the following relative rates (k_obs) [Cu^I(pmea)]⁺ > [Cu^I(pmap)]⁺ > [Cu^I(tmpa)(thf)]⁺. However, the redox potentials as related to ΔG^{{ring operator}} for the reaction, [E_1/2(oxidant) - E_1/2(reductant)], predict the trend [Cu^I(tmpa)(CH₃CN)]⁺ > [Cu^I(pmea)]⁺ > [Cu^I(pmap)]⁺. A detailed description of Cu^I-to-Cu^II structural changes is provided and these likely influence the observed [(L)Cu^I]⁺/(F₈)Fe^III(Cl) reaction rates. Analogous chemistry is described for a heme-copper system utilizing a heterobinucleating ligand (⁶L), which is comprised of a TMPA like moiety tethered at the 6-position of one of the pyridyl donors to a F₈-like heme. Kinetic/mechanistic insights were obtained from transient absorption spectroscopic monitoring in CH₃CN following photoejection of carbon monoxide from [(⁶L)Fe^II(CO)?Cu^II(Cl)].

Full text of DOI:10.1016/j.ica.2007.08.016

Available online at www.sciencedirect.com
Inorganica Chimica Acta 361 (2008) 1100–1115
www.elsevier.com/locate/ica
Intermolecular versus intramolecular electron-/atom- (ClÅ) transfer
in heme-iron and copper pyridylalkylamine complexes
a
a
b
b
H. Christopher Fry , Heather R. Lucas , Lev N. Zakharov , Arnold L. Rheingold ,
a,
a,
*
Gerald J. Meyer *, Kenneth D. Karlin
a
The Johns Hopkins University, Department of Chemistry, 3400 North Charles Street, Baltimore, MD 21218, United States
b
University of Delaware, Crystallography Laboratory, Department of Chemistry, Newark, DE 19716, United States
Received 12 August 2007; accepted 18 August 2007
Available online 25 August 2007
Dedicated to Edward Solomon
Abstract
While outer-sphere electron-transfer reactions come with a firm experimental and theoretical basis, less is known about the redox
reactions occurring by atom-transfer. In the present study, relevant reactions occur upon mixing, (F8)FeIII(Cl) (F8 is tetrakis(2,6-diflu-
orophenyl)porphyrinate) with a series of [(L)CuI]+ complexes (L are tripodal tetradentate pyridylalkylamine ligands varying in effective
chelate ring sizes: TMPA, PMEA, PMAP) to form (F8)FeII and [(L)CuII(Cl)]+. The electron-/atom- (ClÅ) transfer process is characterized
1
by electrochemical measurements as well as UV–Vis, H NMR, and EPR spectroscopies. Stopped-flow UV–Vis spectroscopy in THF
indicate the following relative rates (kobs) [CuI(pmea)]+ > [CuI(pmap)]+ > [CuI(tmpa)(thf)]+. However, the redox potentials as related
to DGꢀ for the reaction, [E1/2(oxidant) ꢁ E1/2(reductant)], predict the trend [CuI(tmpa)(CH3CN)]+ > [CuI(pmea)]+ > [CuI(pmap)]+. A
detailed description of CuI-to-CuII structural changes is provided and these likely influence the observed [(L)CuI]+/(F8)FeIII(Cl) reaction
rates. Analogous chemistry is described for a heme–copper system utilizing a heterobinucleating ligand (6L), which is comprised of a
TMPA like moiety tethered at the 6-position of one of the pyridyl donors to a F8-like heme. Kinetic/mechanistic insights were obtained
from transient absorption spectroscopic monitoring in CH3CN following photoejection of carbon monoxide from [(6L)FeII(CO)ꢂ ꢂ ꢂ
CuII(Cl)].
ꢀ 2007 Elsevier B.V. All rights reserved.
Keywords: Copper; Heme; Redox; Electron-/atom-transfer
1. Introduction
1) first leads to heme-superoxo transients (e.g.,
FeIII–ðO2ꢁÞ ꢂ ꢂ ꢂ CuI species) which then react further to
form peroxo-bridged FeIII–ðl-g2 :g1-O22ꢁÞ–CuII complexes
[1–3]. The latter have been characterized in detail, in con-
nection with modeling the active site chemistry of heme–
copper oxidases which reduce O2 to water [1,4,5]. Reduced
complexes are also known to bind carbon monoxide;
[(6L)FeIIꢂ ꢂ ꢂCuI]+ (Scheme 1) ligates CO at one or both
metal centers. The heme mono-carbonyl, [(6L)FeII(CO)ꢂ ꢂ ꢂ
CuI]+, has proven to be a useful tool in studying photoin-
duced reaction dynamics [6–8].
In the last few years, we have focused considerable
attention on the O2 and CO chemistry of hemes together
with copper ion complexes [1]. Some specific complexes
are shown in Scheme 1, where either as individual compo-
nents, i.e., 1:1 mixtures of [CuI(tmpa)(CH3CN)]+ tmpa2)
(
and (F8)FeII (1), or by using the binucleating ligand com-
plex [6LFeIIꢂ ꢂ ꢂCuI]+, we have shown that reaction of O2
with these reduced FeIIꢂ ꢂ ꢂCuI mixtures/assemblies (Scheme
In earlier work [9], we reported what turned out to be a
useful alternative synthesis of (F8)FeII (1) via introduction
of [CuI(tmpa)(CH3CN)]+ (tmpa2) to [(F8)FeIII(Cl)] (10-Cl) in
*
Corresponding authors.
E-mail address: karlin@jhu.edu (K.D. Karlin).
0020-1693/$ - see front matter ꢀ 2007 Elsevier B.V. All rights reserved.
doi:10.1016/j.ica.2007.08.016

H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
1101
In addition to [CuI(tmpa)(CH3CN)]+
(
tmpa2), a related
series of pyridylalkylamine tripodal tetradentate ligand–
copper complexes (R2), [CuI(pmea)]+, [CuI(pmap)]+, and
[CuI(tepa)]+ (Scheme 2), have been employed. This allows
for a systematic examination of how changes in the nature
of the ligand influence electron-/atom-transfer chemistry
while employing the same oxidant, (F8)FeIII(Cl) (10-Cl). It
has been presented elsewhere that systematically expanding
the five-membered chelate rings by one carbon atom, such
as in the ligand series TMPA ! PMEA ! PMAP !
TEPA, significantly changes the coordination geometry
(based on X-ray diffraction studies) around the copper(II)
ion center in complexes derived from these, while there
are also significant changes in CuII/I redox potentials [10–
12]. Therefore, study of these systems affords the ability
to observe the process of electron-/atom- (ClÅ) transfer as
a function of copper coordination geometry and redox
properties.
Scheme 1.
deaerated toluene (Scheme 2). On the basis of classical
inorganic chemistry redox reaction considerations, it seems
likely that this process occurs by an inner-sphere electron-
transfer mechanism involving a bridging chloride, i.e.,
there is formally a net transfer of a chlorine atom (ClÅ) from
heme to copper. Here, this fundamentally interesting pro-
cess is studied further by examination of related chemical
systems which involve both intra- and inter-molecular elec-
tron-/atom- (ClÅ) transfer reactions.
6
The binucleating ligand L was also employed here in
order to provide further insight into the electron-/atom-
(ClÅ) transfer reactions and for comparison of intra- versus
inter-molecular reaction chemistry. Our previous success in
utilizing 6L to provide the framework necessary to stabilize
bridging O2 intermediates as FeIII–ðO22ꢁÞ–CuII complexes
(vide supra) [3,13] appears ideal for probing the reactions
Scheme 2.

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H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
involving an intramolecular chlorine atom transfer with
possible bridging Clꢁ intermediate. For this purpose, we
have in fact been able to generate [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+
(3FeCl) or species derived from it (as outlined in Section
3) and can indeed observe intramolecular redox-atom
transfer.
2.2. Syntheses of complexes
The following ligand–metal complexes have been previ-
ously synthesized and characterized: [CuI(CH3CN)4]-
B(C6F5)4 [22], [CuI(tmpa)(CH3CN)]B(C6F5)4 tmpa2) [15],
(
[CuI(pmea)]B(C6F5)4
(
pmea2) [16], [CuI(pmap)]B(C6F5)4
Various analytical techniques including 1H NMR, EPR,
electrochemistry, and UV–Vis spectroscopy have been
employed to characterize this electron-/atom- (ClÅ) transfer
process as a function of copper–ligand environment. In
addition, the relative rates of this process have been exam-
ined by transient absorption spectroscopy. In summary, we
report how changes in the nature of the ligand influence
LCuI/LCuII geometries, redox properties, the driving force
(DGꢀ) and rate of electron-/atom- (ClÅ) transfer from heme
to copper. We also report a novel means by which ClÅ atom
transfer reactions can be photoinitiated.
(
pmap2) [16], (F8)FeIII(ClÅ) (10-Cl) [17] and (6L)FeII [3].
The ligands TEPA [11] and 6L [13] was synthesized as pre-
viously described.
2.2.1. Synthesis of [CuI(tepa)]B(C6F5)4 (tepa2)
[CuI(CH3CN)4]B(C6F5)4 (300 mg, 0.33 mmol) was
placed in a 100 mL Schlenk flask equipped with a stirbar.
To this solid, TEPA (110 mg, 0.33 mmol) dissolved in dried
and deaerated diethyl ether was added under an Ar
atmosphere. The resulting yellow solution was stirred for
30 min. Dried and deaerted heptane (100 mL) was added
to precipitate the complex. After stirring for 1 h, the
supernatant was decanted. The remaining yellow oily solid
was dissolved again in dried and dearated diethyl ether
(10 mL) and precipitated with deaerated heptane (100
mL). The resulting solid was placed in vacuo to yield
295 mg (83%) of a yellow powder. Elemental Analysis: Cal-
culated: C, 50.28; H, 2.25; N, 5.21. Experimental: C, 50.11;
H, 1.97; 5.14.
2. Experimental
2.1. Materials and methods
All reagents were purchased from commercial sources.
All air sensitive compounds were synthesized utilizing
Schlenk techniques. Tetrahydrofuran solvent was purified
using MBraun alumina solvent purification columns. Ace-
tonitrile was dried and distilled over calcium hydride
under an argon atmosphere. The air-sensitive compounds
and dried, degassed (purging with argon 1 h) solvents were
introduced into an MBraun Labmaster 130 dry box. UV–
Vis absorption spectra were obtained using a Hewlett–
Packard 8452A or a Hewlett–Packard 8453 diode array
spectrophotometer. Low temperature measurements (ꢁ80
to 0 ꢁC) were made using a Neslab ULT-95 controller
by circulating methanol through a two window optical
dewar as previously described [3,14]. UV–Vis spectro-
scopic samples of (F8)FeIII(Cl) in dried, air-free CH3CN
were added to samples of [(L)CuI]B(C6F5)4 in equimolar
concentrations in the dry box and transferred to a
10 mm Quark Glassware Schlenk quartz cuvette equipped
with a 14/20 rubber septum. X-band electron paramag-
netic resonance (EPR) spectroscopic studies were per-
formed using a Bruker EMX magnet operating at a
frequency of 9.4 GHz; 1–2 mM samples were prepared in
50:50 CH3CN:toluene and frozen in liquid N2 before cool-
ing further to 12–16 K for spectroscopic measurements.
6
2.2.2. Synthesis of LFeIII(Cl) Æ 2H2O Æ CH2Cl2
6
The ligand L (900 mg) was dissolved in dimethylform-
amide (150 mL) under Ar and placed in a 250 mL, three-
neck round bottom flask equipped with a stirbar and
condenser. This was then attached to a Schlenk line. The
solution was purged with argon for 30 min and was then
heated until the flask was warm to the touch. At this time,
FeIICl2 Æ 4H2O was added to the solution under positive
argon flow. After 2.5 h, the reaction was expose to air
and was set to cool to room temperature. The DMF was
removed by rotary evaporation. The crude product was
dissolved in methylene chloride and was then added to a
stirring solution of saturated NaOH (700 mL). This mix-
ture was then stirred overnight. The methylene chloride
layer was then dried over sodium sulfate and adsorbed
onto alumina. Purification by column chromatography
(1:1 hexanes/ethyl acetate) yielded(6L)FeIII(OH). This
product was then dissolved in 100 mLCH2Cl2 and trimeth-
ylsilylchloride (TMS-Cl, 1.0 mL) was added. After 3 h, the
volatiles were removed by rotary evaporation, yielding a
purple product. UV–Vis (CH3CN): Soret, 412 nm; a, 504,
572, 630 nm. MALDI-TOF MS: 1131.0 m/z. Elemental
Analysis: Calculated: C, 61.43; H, 3.71; N, 8.96. Experi-
mental: C, 61.61; H, 3.70; N, 8.34.
NMR spectroscopy was performed using
a Varian
400 MHz instrument. All spectra were recorded at room
temperature in 5 mm o.d. NMR tubes. Reported spectra
were generated using MestReC NMR processing software.
Matrix-assisted laser desorption ionization time-of-flight
(MALDI TOF) mass spectra were recorded with a Kratos
2.2.3. Synthesis of [(6L)Feꢂ ꢂ ꢂ(Cl)ꢂ ꢂ ꢂCu]B(C6F5)4 (3)
Analytical Kompact MALDI
4
mass spectrometer
In a 150 mL Schlenk flask, (6L)FeIII(Cl) (100 mg,
equipped with a 337 nm nitrogen laser (20 kV extraction
voltage). No matrix was necessary for the MALDI-TOF
samples as the presence of porphyrin ligand itself proved
to be sufficient for laser desorption.
0.080 mmol)
and
[CuI(CH3CN)4]B(C6F5)4
(78 mg,
0.086 mmol) were dissolved in 10 mL of tetrahydrofuran.
The solution was stirred for 1 h. During this time,
100 mL of heptane (dried and distilled) was bubbled with

H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
1103
argon and this was added to the stirring solution which
resulted in precipitation of a purple solid. The heptane
was removed via cannula and the resulting purple solid
was dried in vacuo. UV–Vis (CH3CN, ꢁ40 ꢁC) Soret,
424 nm; a, 528 nm; (THF, ꢁ40 ꢁC) Soret, 424 nm; a,
541 nm. MALDI-TOF MS: 1193.0 m/z. Elemental Analy-
sis: Calculated: C, 54.76; H, 2.11; N, 5.87. Experimental:
C, 54.37; H, 2.28; N, 5.52.
ture variation from 183 to 298 K. The stopped-flow
apparatus was coupled to a J&M Tidas diode-array spec-
trophotometer equipped with J&M Tidas Spectralys 200
software. All kinetic processes were analyzed using the
first-order integrated rate law with ORIGIN 6.1 software.
Samples were prepared in the dry box and transferred to
an additional funnel designed to attach to the stopped-flow
apparatus in order to load the syringes under an inert
atmosphere. The apparatus used was similar to that previ-
ously employed and it was used in a like manner [20].
2.3. Electrochemistry
Cyclic voltammetry was carried out using a Bioanalyti-
cal Systems BAS-100B electrochemistry analyzer. The cell
consisted of a 25 mL three-necked round bottom flask
equipped with rubber septa. Each neck possessed one of
the following: a glassy carbon working electrode, a satu-
rated calomel reference electrode equipped with a salt
bridge necessary for organic solvents or a Ag/AgPF6 elec-
trode in CH3CN, and a platinum wire as the auxiliary elec-
trode. Ferrocene was used as an external standard. The
measurements were performed at room temperature under
an argon atmosphere in CH3CN solutions, containing
100 mM tetrabutylammonium hexafluorophosphate elec-
trolyte, using 1–3 mM complex concentrations.
2.6. X-ray diffraction
X-ray quality crystals of [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuII(Cl)]-
B(C6F5)4 (4) were obtained from dissolving [(6L)Feꢂ ꢂ ꢂ
(Cl)ꢂ ꢂ ꢂCu]B(C6F5)4 (3) in MeOH/CH2Cl2 (<1 mL) and
layering with heptane (5 mL) in a 5 mm o.d. NMR tube
under Ar. After 4-days purple crystals were obtained.
The X-ray data was obtained and analyzed at the Univer-
sity of Delaware (laboratory of A.L. Rheingold). The
structure was solved using direct methods, completed by
subsequent difference, Fourier syntheses, and refined by
full-matrix least-squares procedures. All non-hydrogen
atoms were refined with anisotropic displacement coeffi-
cients, and the hydrogen atoms were treated as idealized
contributions. All software and sources of the scattering
factors are contained in the SHELXTL (5.1) program library
(G. Sheldrick, Siemens XRD, Madison, WI). Solvent
molecules, CH2Cl2 and CH3OH, in the crystal structure
were highly disordered and were treated by SQUEEZE
[21]. Correction of the X-ray data by SQUEEZE, 210 elec-
tron/cell, was close to the required values 204 electron/cell
for four CH2Cl2 and two CH3OH in the full unit cell.
2.4. Transient absorption spectroscopy
Transient absorbance measurements were performed
using an Nd:YAG (Continuum Surelite III) laser on a pre-
viously described apparatus [18]. Samples of [(6L)FeII-
(CO)ꢂ ꢂ ꢂCuII(Cl)]B(C6F5)4 (5) were prepared in situ upon
dissolution of [(6L)Feꢂ ꢂ ꢂ(Cl)ꢂ ꢂ ꢂCu]B(C6F5)4 (3) in acetoni-
trile and saturating with CO(g). The samples were irradi-
ated with 532 nm pulsed light at 2–10 mJ/cm2 and
protected from the probe beam with appropriate UV fil-
ters. For full transient absorption spectra, 0.5–3 lM sam-
ples were used in a 1.0 cm pathlength cuvette and kinetic
traces were monitored from 320 to 650 nm with varying
steps (every 5–10 nm) averaging 40 transients per
wavelength.
The CO concentration was varied in all solvents studied
in order to determine bimolecular rate constants, kCO
(Mꢁ1 sꢁ1). CO was varied (1–10 mM in 1 mM steps) as pre-
viously described using a gas mixer (N2 was used as the
supporting inert gas) [19]. Four wavelengths were moni-
tored in the Soret region averaging 60–120 transients per
wavelength and the observed rate constants were averaged
at each CO concentration studied. Plotting the observed
rate constants, kobs versus CO concentration resulted in a
linear fit. The slope corresponds to the bimolecular rate
of recombination, kCO (Mꢁ1 sꢁ1).
3. Results and discussion
3.1. Synthesis
The syntheses of the mononuclear copper(I) species (R2),
[CuI(tmpa)(CH3CN)]B(C6F5)4
(
tmpa2),
[CuI(pmea)]-
(
pmap2), are
B(C6F5)4
(
pmea2), [CuI(pmap)]B(C6F5)4
described in a separate publication [16]; that for [CuI-
(tepa)]B(C6F5)4 (tepa2) (Scheme 2) is given in the Experi-
mental Section. The cupric chloride species (R20-Cl),
[CuII(tmpa)(Cl)]B(C6F5)4
(
tmpa20-Cl), [CuII(pmea)(Cl)]B-
(C6F5)4 (pmea20-Cl), and [CuII(pmap)(Cl)]B(C6F5)4 (pmap20-
Cl), were produced in situ upon addition of the respective
cuprous salt to the heme–chloride complex 10-Cl in aceto-
nitrile (CH3CN) or tetrahydrofuran (THF) solvent under
argon (Scheme 2). With the presence of such strong heme
chromophores, a typical color change from brown to
bright red is visible on the bench-top upon reduction of
the ferric tetraarylporphyrins to their ferrous counterparts.
Procedures for metallation of the free base porphyrins,
F8 and 6L, with FeIICl2 Æ 4H2O to produce the correspond-
ing (F8)FeIIICl (10-Cl) [17] (Scheme 2) and (6L)FeIIIOH [3]
ferric compounds were previously published; these com-
2.5. Stopped-flow spectroscopy
Time-resolved UV–Vis spectroscopy was performed
using a Hi-Tech Scientific SF-40 stopped-flow apparatus
equipped with a low temperature dewar allowing tempera-

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H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
plexes are already fully characterized. The target ferric
chloride species of the heterobinucleating ligand 6L,
(6L)FeIII(Cl) with empty TMPA moiety, was produced
upon washing (6L)FeIII(OH) with a mild acid, trimethylsi-
lylchloride (TMS-Cl); treatment with a stronger acid, such
as HCl, results in cleavage of the heterobinucleating ligand
ether linkage.
cially as concerns the copper ion environment. The heme
portion of 4 follows what is typically expected of five-coor-
dinate
high-spin
(porphyrinate)FeIII–X
complexes
[17,23,24], with 4 having an average Fe–N distance of
˚
˚
2.06 A and with the iron ion positioned 0.46 A out of the
heme plane. The copper moiety of 4 exhibits a distorted
square pyramidal (sp) geometry (also see discussion
below). A distinctive aspect of the present structure is that
the tethered pyridyl moiety containing the cupric ion
swings away from the heme, possibly as a result of the
two chloride anions repelling one another or perhaps as a
preferential crystal and crystal packing arrangement.
Introduction of [CuI(CH3CN)4]B(C6F5)4 [22] to metal-
late the non-heme moiety of (6L)FeIII(Cl) in deaerated
THF solvent also results in chloride transfer as can be
observed macroscopically. However, the electron-/atom-
(ClÅ) transfer from iron to copper is incomplete leading to
a
mixture of [(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+ (3CuCl
)
and
[(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl); this electron-/atom- (ClÅ)
transfer phenomena will be described throughout in detail.
The equilibrium position (Scheme 3), 3CuCl versus 3FeCl, is
dependent on solvent interaction and concentration; there-
fore, the isolated solid and apparent mixture is from here
on referred to as [(6L)Feꢂ ꢂ ꢂ(Cl)ꢂ ꢂ ꢂCu]+ (3).
3.3. Relevant structural aspects of copper(I/II) complexes
Geometric variation of the ligand environment around
the copper ion like presented here is of fundamental inter-
est for a further understanding of the various factors
involved when converting copper(I) to copper(II) via elec-
tron-/atom- (ClÅ) transfer reactions. In addition, the geom-
etry change which results upon reduction of (F8)FeIII(Cl)
(10-Cl) to (F8)FeII (1) contributes to this process, but we
employed 10-Cl as our single oxidizing species to minimize
variables that may alter our understanding of these
reactions.
3.2. X-ray crystallography of [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuII(Cl)]+
(4)
The attempted crystallization of [(6L)Feꢂ ꢂ ꢂClꢂ ꢂ ꢂCu]+ (3)
in a MeOH/CH2Cl2/heptane solvent mixture led to crystals
of a disproportionation product, [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuII(Cl)]+
(4), which was characterized by X-ray crystallography,
Fig. 1. This compound is relevant in order to deduce struc-
tural aspects of the heme–Cu assembly being studied, espe-
When copper is oxidized from copper(I) to copper(II) in
typical non-protein ligand environments, a drastic geomet-
ric change from the ideal four coordinate tetrahedral
cuprous species (3d10) to a five coordinate square-based
pyramidal (sp) or trigonal bipyramidal (tbp) cupric (3d9)
species most often occurs. For [CuI(tmpa)(CH3CN)]+
1
(
tmpa2), H NMR investigations have indicated an internal
fluxional process where dissociation and re-association of a
pyridyl arm occurs rapidly. The tetradentate ligand is
essentially coordinated in an N3 fashion with an additional
solvent ligand present (Scheme 4) [20,25,26]. Thus, the
structural rearrangement necessary to convert tmpa2 to the
perfect trigonal bipyramidal tmpa20-Cl complex is quite
drastic (Scheme 4). On the contrary, much less dramatic
geometric changes occur in the cases of [CuI(pmea)]+
(
pmea2) and [CuI(pmap)]+ (pmap2), as all of the pyridyl nitro-
gens are coordinated to the cuprous metal center in solu-
tion [11]. The addition of a chloride anion by 10-Cl to
pmea2 and pmap2 expands the ligand–metal coordination
number from four to five, in forming their corresponding
cupric analogues, pmea20-Cl and pmap20-Cl, as there is not
Scheme 3.
Fig. 1. ORTEP representation and drawing based on the X-ray structure
of [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuII(Cl)]+ (4). See text for further discussion.
Scheme 4.

H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
1105
a solvent molecule coordinated in the CuI ion complex
[11,16].
Based on the structural index parameter s, defined by
Reedijk and Addison [27] to describe pentacoordinate com-
plexes, a perfect tbp geometry has a s value of 1, like
[CuII(tmpa)(Cl)]+, and sp coordination has a s value of 0,
similar to that of [CuII(tepa)(Cl)]+, Scheme 5 [28]. The
TEPA framework of [CuII(tepa)(Cl)]+ houses a geometry
closest to the ideal square planar geometry desired by 3d9
cupric systems. The sp cupric complexes pmea20-Cl and
pmap20-Cl have s values of 0.12 and 0.08, respectively [11].
Similarly, the s value for the copper(II)–chloride moiety
of the heme–copper complex [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuII(Cl)]+
(4) is 0.19, where the axial ligand is the pyridyl group com-
prising the linker between the heme and copper-chelate in
Fig. 2. UV–Vis spectra of the a-region for (F8)FeIII(Cl) (10-Cl) in CH3CN
(black line) and the reduced (F8)FeII (1) complex upon addition of
[(L)CuI]+ to 10-Cl in CH3CN (red line); where L is TMPA, PMEA,
PMAP.
6
˚
L; Cu–N = 2.32 A. This s value is drastically different
than the mononuclear analogue tmpa20-Cl. However, we
previously established that cupric chloride TMPA analogs
possessing 60-substituents on one of the pyridyl groups
begin to distort their coordination geometry from perfect
tbp to sp [29]. The square-based pyramidal copper geome-
try of the copper(II)–chloride moiety of [(6L)FeIIꢂ ꢂ ꢂ
CuII(Cl)]+ (3CuCl) is assumed based on similarities to 4
and is therefore most similar in structure to pmea20-Cl. Such
large structural differences amongst the various mononu-
clear copper(II)–chloride species (R20-Cl) may influence
the electron-/atom-transfer (ClÅ) chemistry, vide infra.
bances at 424 nm (Soret) and 525 nm (a) [17], Fig. 2. No
evidence for a reaction was observed upon addition of
[CuI(tepa)]+ tepa2) to 10-Cl in CH3CN (Scheme 2).
(
Dissolution of the heterobinucleating system, [(6L)Feꢂ ꢂ ꢂ
(Cl)ꢂ ꢂ ꢂCu]+ (3) does not appear to result in ClÅ transfer
from heme chemistry in CH3CN at 298 K, as indicated
by the a-region possessing the three bands typical of the
ferric chloride species, Fig. 3. However, at lower concentra-
tions of 3, the UV–Vis features (p–p*) broaden, the Soret
absorbance shifts to lower energy (420 nm), and the LMCT
band (380 nm) decreases in intensity suggesting partial ClÅ
transfer from heme to copper, Fig. 4, i.e., there is an
equilibrium mixture of [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl) and
[(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+ (3CuCl), Scheme 3, also see further
discussion below.
3.4. Spectroscopic characterization of the electron-/atom-
(ClÅ) transfer reactions
3.4.1. UV–Vis spectroscopy
Iron porphyrinates have distinctive spectroscopic UV–
Vis signatures, which dominate the charge transfer and/
or d–d absorptions of copper(I) and/or copper(II) com-
plexes. In CH3CN at 298 K, (F8)FeIII(Cl) (10-Cl) possesses
a distinctive ligand-to-metal charge transfer (LMCT) band
at 380 nm [30] along with p–p* features observed within the
Soret region at 414 nm and three less intense absorbances
in the a-region at 504, 573, and 620 nm [17], Fig. 2. Addi-
Nitrile solvents such as CH3CN strongly coordinate to
CuI and can even inhibit CuI/O2 chemistry, as has been
observed in tridentate ligand–copper/heme systems pos-
sessing similar pyridyl-ethyl-amine donor moieties [14,31].
tion of [CuI(tmpa)(CH3CN)]+
(
tmpa2), [CuI(pmea)]+
(
pmap2) to 10-Cl in CH3CN results
(
pmea2), or [CuI(pmap)]+
in electron-/atom- (ClÅ) transfer reactions leading to com-
plete formation of [LCuII(Cl)]+ (R20-Cl) and (F8)FeII (1),
Scheme 2, the latter of which has characteristic absor-
Fig. 3. UV–Vis spectra of the a-region indicate the presence of the
[(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+:[(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+ (3FeCl):(3CuCl) equilibrium in
THF (black, clearly two peaks) and the predominantly
[(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl) equilibrium species in CH3CN (red) at room
temperature.
Scheme 5.

1106
H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
Fig. 5. UV–Vis spectra at 233 K in THF of [(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+ (3CuCl
)
Fig.
4. UV–Vis
spectra
indicating
the
presence
of
the
þ
III
II
[(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+:[(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+ (3FeCl): (3CuCl) equilibrium in
THF (black) and the predominantly [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl) in
CH3CN (red) at room temperature.
6
(in black) and ½ð LÞFe ðO2ꢁÞ ꢂ ꢂ ꢂ Cu ðClÞꢃ (in red).
Therefore, ClÅ transfer within 3 to copper is most likely
hindered by CH3CN solvent coordination to the CuI ion,
see Scheme 6a. In an attempt to make the CuI ion of
[(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl) more accessible for efficient
electron-/atom-transfer, the experiments were repeated in
THF as solvent. However, in THF at all concentrations,
there is a mixture of [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl) and
[(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+ (3CuCl), as indicated by the broad
absorbances observed in the Soret and a-region, Figs. 3
and 4. In contrast to the weak non-existent coordination
of THF to CuI ions, ferrous hemes bind THF to form more
stable complexes [8,17]. Therefore, such inefficient
ClÅ transfer may be due to stabilization of the ferric chlo-
ride moiety by coordination of an axial base (THF), see
Scheme 6b.
By decreasing the temperature of the solution of 3 to
233 K in either CH3CN and THF at low complex concen-
tration, the equilibrium shifts towards [(6L)FeIIꢂ ꢂ ꢂ
CuII(Cl)]+ (3CuCl), as confirmed by the sharpening of the
Soret- and a-absorbances as well as the decrease in LMCT
intensity (Fig. 5). Interestingly, upon addition of dioxygen,
þ
III
II
6
a stable superoxide species ½ð LÞFe ðO2ꢁÞ ꢂ ꢂ ꢂ Cu ðClÞꢃ is
generated in both CH3CN and THF, Soret = 414 nm,
a = 538 nm (Scheme 6c and Fig. 5). In the absence of chlo-
ride ligand, the reaction product of [(6L)FeIIꢂ ꢂ ꢂCuI]+ and
dioxygen is established as a bridging peroxo species,
þ
III
II
½ð LÞFe ðO22ꢁÞCu ꢃ [3]. Here, in the absence of an elec-
tron from copper(I) to further reduce dioxygen to a peroxo
species, we observe a stable heme-superoxide species; the
UV–Vis absorptions are very characteristic and the same
6
Scheme 6.

H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
1107
heme-superoxo formation chemistry has been established
for reaction of O2 with (F8)FeII and (6L)FeII (without
CuI ion in tethered moiety) which form ½ðF8ÞFeIIIðO ꢁÞꢃþ
2
þ
III
6
and ½ð LÞFe ðO2ꢁÞꢃ , respectively [17,32–34].
1
3.4.2. H NMR spectroscopy
Multiple spin-state configurations are available for FeII/
FeIII coordination compounds leading to differences in
reactivity amongst species of the same redox state. In
1
CD3CN at room temperature, the H NMR spectrum of
(F8)FeIII(Cl) (10-Cl) confirms a high spin 3d5 species as
indicated by the paramagnetically shifted peak at d
80 ppm for the pyrrole hydrogens [17]. Upon addition of
[CuI(tmpa)(CH3CN)]+
[CuI(pmap)]+
( (
tmpa2), [CuI(pmea)]+ pmea2), or
(
pmap2), the pyrrole-H resonances are shifted
upfield to d 25 ppm (Fig. 6), representative of an equilib-
rium mixture of high- and low-spin (F8)FeII (1), Fig. 6.
In previous work, the magnetic moment of 1 was deter-
mined at variable temperature via the Evans NMR method
(l = 4.1) [17]. Upon lowering the temperature to 233 K,
the pyrrole hydrogens displayed anti-Curie behavior, shift-
ing upfield to d 10 ppm in CH3CN, indicative of full forma-
tion of low-spin six-coordinate heme (two axially bound
CH3CN such as in Scheme 6d) [17]. The absence of the fer-
ric pyrrole-H resonances at d 80 ppm upon adding [LCuI]+
(R2) to 10-Cl suggests that electron-/atom-transfer is not
affected by the spin state of the ferrous reaction product
and that full conversion occurs, i.e., ClÅ transfer, forming
[LCuII(Cl)]+ (R20-Cl) and (F8)FeII (1), Scheme 2.
Fig. 7. 1H NMR spectrum of [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl) in CD3CN.
The downfield paramagnetically shifted resonances centered at dpyrrole
=
82 ppm are consistent with the presence of only the high spin ferric
chloride porphyrin complex.
high spin species as revealed by the [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+
(3FeCl) pyrrole peak at d 80 ppm and the pyrrole resonances
representative of [(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+ (3CuCl) at 50–
55 ppm. Integration of the broad pyrrole resonances results
in a roughly 1:1 mixture of the two species. Upon introduc-
tion of 1,5-dicyclohexylimidazole (DCHIm), a strongly
coordinating axial base for heme, the equilibrium is pushed
towards the [(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+ (3CuCl) conformer. Simi-
larly, repeating the experiment in more strongly coordinat-
ing solvent, pyridine-d5, results in a shift in equilibrium
towards [(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+ (3CuCl), as observed by the
presence of a pyrrole signal at d 8.8 ppm, consistent with
a low-spin ferrous species [(6L)FeII(py)2ꢂ ꢂ ꢂCuII(Cl)]+ like
that previous observed for [(F8)FeII(py)2], Fig. 8 [17].
Therefore, the presence of an organic base like THF,
1
The H NMR spectrum of [(6L)Feꢂ ꢂ ꢂClꢂ ꢂ ꢂCu]+ (3) in
CD3CN (Fig. 7) looks nearly identical to the empty tether
complex, (6L)FeIII(Cl), suggesting that the ClÅ remains on
the heme, i.e., no electron-/atom-transfer has taken place,
as supported by UV–Vis spectroscopic analysis (vide
supra). When changing to the more coordinating solvent
for hemes, THF-d8, we observe a mixture of FeII and FeIII
Fig. 6. 1H NMR spectrum following the reaction of (F8)FeIII(Cl) (10-Cl)
with [CuI(tmpa)(CH3CN)]+
(
tmpa2) in CD3CN, d (pyrrole) = 25 ppm. The
Fig. 8. 1H NMR spectrum of [(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+ (3CuCl) in pyridine-d5.
No paramagnetic resonances were observed as the coordinating ability of
pyridine forces the equilibrium to the low-spin d6 ferrous porphyrin
complex. dpyrrole = 8.8 ppm.
product is a mixture of high (with one ligated axial CD3CN) and low-spin
(two ligated axial CD3CN molecules) (F8)FeII (1) plus [CuII(tmpa)(Cl)]+
(
tmpa20-Cl).

1108
H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
DCHIm, or pyridine-d5 will clearly push the equilibrium
(Scheme 3) entirely to the stabilized low-spin d6 ferrous
state to form [(6L)FeII(Base)2ꢂ ꢂ ꢂCuII(Cl)]+ (Scheme 6d).
The sensitive nature of this reaction in non-coordinating
solvents suggests that the redox potentials of the two metal
centers are nearly identical under such conditions. This
supposition is substantiated by the electrochemical studies,
see below.
3.4.3. Electron paramagnetic resonance (EPR) spectroscopy
X-band EPR analyses were conducted on frozen sam-
ples (12–16 K; 1–2 mM) in CH3CN:toluene (1:1) solvent
mixtures. (F8)FeIII(Cl) (10-Cl) has a typical high spin 3d5
heme signal with g^ = 5.92 and gk ¼ 2:00, Fig. 9. Upon
addition of DCHIm, the signal changes to a low-spin 3d5
species due to coordination of a strong axial base to the
heme, forming (F8)FeIII(Cl)(DCHIm). Instead, addition
of 1 equiv. of [CuI(tmpa)(CH3CN)]+ (tmpa2) to 10-Cl results
in complete silencing of the heme species indicative of the
formation of an FeII 3d6 system and the presence of a
Fig. 10. X-band EPR spectra of [(6L)Feꢂ ꢂ ꢂClꢂ ꢂ ꢂCu]+ (3) (black) and
[(6L)FeII(DCHIm)2ꢂ ꢂ ꢂCuII(Cl)]+ (Scheme 6d) (red).
pyridyl 60-substituent in accordance with previous studies
[29].
reverse axial signal (g^ = 2.01, gk ¼ 2:18, A^ = 96cmꢁ1
)
typical of a trigonal bipyramidal CuII 3d9 system, as previ-
ously established for tmpa20-Cl (Fig. 9) [28,35]. The com-
plete silencing of the iron is consistent with a quantitative
efficient electron-/atom-transfer (ClÅ) process, Scheme 2.
As expected, the EPR spectrum of binuclear
[(6L)Feꢂ ꢂ ꢂClꢂ ꢂ ꢂCu]+ (3) upon dissolution in an CH3CN:tol-
uene (1:1) solvent mixture displays only high spin ferric
chloride; cuprous ion is EPR silent, Fig. 10. Addition of
DCHIm, results in complete silencing of the ferric EPR sig-
nal and appearance of an axial signal (g ꢄ 2) due to the
presence of a copper(II) species, Fig. 10. As opposed to
3.4.4. Infrared spectroscopy
Since [(6L)FeII(Base)1–2ꢂ ꢂ ꢂCuII(Cl)]+ is favored in the
presence of an axial base (i.e., DCHIm, THF, and pyri-
dine-d5), IR spectroscopy was employed to further confirm
complete ClÅ transfer by addition of CO to [(6L)Feꢂ ꢂ ꢂ
Clꢂ ꢂ ꢂCu]+ (3). CO binds to both iron and copper ions in
the reduced state [7]; therefore the location of the coordi-
nated chlorine can be assumed based on the carbonyl
species formed. Previously, the bis-carbonyl complex
[(6L)FeII(CO)ꢂ ꢂ ꢂCuI(CO)]+
addition of CO to 3 in CH3CN, only mCO(Fe) = 1975 cmꢁ1
is observed, indicating formation of [(6L)Fe(CO)ꢂ ꢂ ꢂ
Cu(Cl)]+ (5) (Scheme 6e) and confirming that the equilib-
rium shifts towards the ferrous state with bound CO.
was
characterized
with
m
CO(Fe) = 1975 cmꢁ1 and mCO(Cu) = 2091 cmꢁ1 [7]. Upon
6
the mononuclear F8 system versus the L system, a low-
spin 3d5 iron(III) species is not observed upon addition
of a coordinating base, implying formation of [(6L)FeII(D-
CHIm)2ꢂ ꢂ ꢂCuII(Cl)]+, Scheme 6d. The poorly resolved
axial (but not ‘‘reverse’’) EPR signal is due to the
[(L)CuII(Cl)]+ moiety, distorted from tbp because of the
3.5. Electrochemical estimates of the driving force (DGꢀ) for
the electron-/atom- (ClÅ) transfer reactions
Heme species can involve many redox changes, over a
number of oxidation states, i.e., IV/III, III/II, and II/I
[24,36–38]. In addition, the porphyrin ring can be oxidized
or reduced to a cation or anion radical [24]. To describe
electron-/atom-transfer chemistry in the present systems,
FeIII/FeII and CuII/CuI redox processes are of primary con-
cern. Cyclic voltammetric measurements were carried out
on CH3CN solutions of each complex under investigation;
E1/2 values are reported in Tables 1 and 2. All voltammo-
grams were classified as quasi-reversible having ipc/ipa ꢅ 1
with half-wave potentials in the range E1/2 = +60 to
ꢁ740 mV versus Fc/Fc+ with DE = 60–300 mV.
In the present study, the redox potentials (E1/2) of the
mononuclear cupric chloride species (R20-Cl) and (F8)FeII
(1) were determined in situ upon dissolution of the comple-
mentary cuprous salt (R2) in an CH3CN solution contain-
ing 1 equiv. of (F8)FeIII(Cl) (10-Cl), (Scheme 2) and
Fig. 9. X-band EPR spectra of (T = 12–16 K) of (F8)FeIII(Cl) (10-Cl)
(black) and the product mixture of 10-Cl and [CuI(tmpa)(CH3CN)]+
(
tmpa2) (red).

H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
1109
Table 1
Redox potentials recorded in CH3CN at room temperature of various synthetic hemes and copper complexes
Complex
E1/2 (mV vs. Fc/Fc+)
DE (mV)
Redox couple
[CuI(tmpa)(CH3CN)]+
(
tmpa2)
ꢁ400
ꢁ740
ꢁ360
ꢁ670
ꢁ250
ꢁ590
60
100
130
300
130
300
110
100
100
80
CuII/I
CuII/I
CuII/I
CuII/I
CuII/I
CuII/I
CuII/I
FeIII/II
FeIII/II
[CuII(tmpa)(Cl)]+ tmpa20-Cl)
(
[CuI(pmea)]+
[CuII(pmea)(Cl)]+
[CuI(pmap)]+ pmap2)
[CuII(pmap)(Cl)]+ pmap20-Cl)
[CuI(tepa)]+ tepa2)
(
pmea2)
(
pmea20-Cl)
(
(
(
[(F8TPP)FeII] (1)
ꢁ140
[(F8TPP)FeIII(Cl)] (10-Cl)
ꢁ530
See text for references to previously measured values.
Table 2
Half-wave potentials measured in CH3CN at room temperature for [(6L)Feꢂ ꢂ ꢂClꢂ ꢂ ꢂCu]+ (3) and derivativesa
Complex
E1/2 (mV vs. Fc/Fc+)
DE (mV)
Redox couple
(6L)FeII
ꢁ170
ꢁ560
ꢁ530
ꢁ490
ꢁ650
100
100
110
60
FeIII/II
(6L)FeIII(Cl)
FeIII/II
[(6L)Feꢂ ꢂ ꢂ(Cl)ꢂ ꢂ ꢂCuI]+ (3)
[(6L)FeII(DCHIm)2ꢂ ꢂ ꢂCuII(Cl)]
[(6L)FeII(DCHIm)2ꢂ ꢂ ꢂCuII(Cl)]b
a
FeIII/II, CuII/I
FeIII/II
CuII/I
110
All measurements were made under 1 atm of argon unless stated otherwise.
Measured under 1 atm of carbon monoxide.
b
(100 mM) tetrabutylammonium hexafluorophosphate elec-
trolyte. Since electron-/atom-transfer does not occur upon
mixing tepa2 and (F8)FeIII(Cl), the E1/2 value of [CuII(te-
pa)(Cl)]B(C6F5)4 in CH3CN was not determined here but
it is known to be nearly 0.5 V more positive than that for
tmpa20-Cl [12].
arm length) of the ligands housing the copper ion, the E1/2
values become more positive, favoring copper(I), see
Fig. 11 and Table 1. The most positive redox potential,
E1/2 = +60 mV, assigned to [CuI(tepa)]+
(
tepa2), is
explained by the near-ideal tetrahedral ligand framework
provided by TEPA which is the most stable geometry for
a cuprous species.
The E1/2 values for the [B(C6F5)4]ꢁ salts of the various
cuprous and cupric chloride species followed the same
trend as was previously obtained for their [LCuI]PF6 and
[LCuII(Cl)]ClO4 analogues in dimethylformamide (DMF)
[11,39]. Upon increasing the chelate ring size (i.e., ligand
Chloride, a strongly coordinating anion, stabilizes the
more highly charged metal ion species, resulting in a nega-
tive shift of the CuII/CuI and FeIII/FeII redox potentials.
The E1/2 values of the cupric chloride species are
ꢄ300 mV more negative than their cuprous partner species,
this value being conserved for each copper species, Table 1.
Such trends in E1/2 values for [LCuII(Cl)]+ (R20-Cl) and
their corresponding [LCuI]+ (R2) compounds are within
the same range and parallel to those from an analogous
study on quinolyl-for-pyridyl systematic ligand variations
[39]. Moreover, the E1/2 value of 1 shifts from ꢁ140 mV
to ꢁ530 mV for 10-Cl in CH3CN, Table 1.
From the redox potentials obtained for the heme and
copper complexes, the relative driving force (DGꢀ) for the
overall electron-/atom- (ClÅ) transfer process was estimated
by calculating [E1/2(oxidant) ꢁ E1/2(reductant)]. The driv-
ing force was found through systematic variation of E1/
2(reductant) as a function of the ligand environment
around copper. The ferric oxidant 10-Cl reacts with all of
the [(L)CuI]+ complexes discussed here, Scheme 2, with
Fig. 11. Cyclic voltammograms of [CuII(tmpa)(Cl)]+
(
tmpa20-Cl) (black),
pmap20-Cl)
the exception of [CuI(tepa)]+ tepa2). The overall negative
(
[CuII(pmea)(Cl)]+ pmea20-Cl) (red), and [CuII(pmap)(Cl)]+
(
(
driving force for the reaction with tepa2 (DEꢁ = ꢁ200 mV)
indicates that a reaction should not occur which is consis-
tent with the UV–Vis results (vide infra). In addition, we
find 10-Cl should oxidize tmpa2 the most effectively, with a
(blue) in CH3CN with (Bu)4NPF6 referenced against Ag/AgPF6. Each
complex was generated in situ from the addition of (F8)FeIII(Cl) (10-Cl) to
[(L)CuI]B(C6F5)4 (R2). The redox wave for the reduced (F8)FeII (1)
complex occurs at ꢁ140 mV (not shown/off-scale) vs. Fc+/Fc.

1110
H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
driving force of DEꢀ = +210 mV while the pmap2 reaction
with 10-Cl should be the least effective, having a much smal-
ler driving force, DEꢀ = +60 mV. When calculating the
driving force values, we were consistent in using those
E1/2 values determined in the presence of a chloride anion
as this has shown to significantly influence the FeIII/II and
CuII/I reduction potentials.
is rendered irreversible yet the CuII/I redox potential does
not change, Table 2. In the case of [(6L)Feꢂ ꢂ ꢂClꢂ ꢂ ꢂCu] (3),
it is interesting to find that upon addition of Lewis bases
known to stabilize ferrous hemes that the redox waves also
separate, meaning the E1/2 values of the two metal ions are
no longer equivalent and the equilibrium lies towards
3CuCl
.
Consistent with the ꢁ390 mV difference in the reduction
potential of (F8)FeIII/II(Cl) and (F8)FeIII/II, the E1/2 value
for (6L)FeIII/II(Cl) shifts from ꢁ560 mV to ꢁ170 mV for
(6L)FeIII/II. Surprisingly, only a small change in the overall
redox potential (E1/2 = ꢁ530 mV) occurs upon addition of
a cuprous salt into the non-heme tethered portion of
(6L)FeIII(Cl) to form [(6L)Feꢂ ꢂ ꢂClꢂ ꢂ ꢂCu] (3), Fig. 12. The
visual presence of only one reversible wave suggests that
the redox waves observed in cyclic voltammetry corre-
sponding to the FeIII/II and CuII/I potentials are overlap-
ping (if not identical). Similarly, overlapping heme/M
(M = Cu, non-heme iron) redox processes have been
observed [40–42]. By considering the Nernst equation, if
E1/2(Cu) = E1/2(Fe) then K ꢅ 1, an equilibrium state
(Scheme 3) is suggested and ꢄ1:1 molar ratios of
[(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl) and [(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+
(3CuCl) are present, as supported by 1H NMR spectroscopy
(vide infra).
3.6. Observed rates of electron-/atom- (ClÅ) transfer as
determined by stopped-flow UV–Vis spectroscopy
Using a stopped-flow kinetics instrument with UV–Vis
spectroscopic monitoring, equimolar THF solutions of
[(L)CuI]+ (R2) and (F8)FeIII(Cl) (10-Cl) complexes were
rapidly mixed at variable low temperatures (188–263 K)
in order to gain further insight into the electron-/atom-
(ClÅ) transfer process, Scheme 2. As discussed, the calculated
relative driving forces predicted relative rates corresponding
to the measured redox potentials of the various cuprous
reductants in the order [CuI(tmpa)(CH3CN)]+
(
tmpa2) >
[CuI(pmea)]+ pmea2) > [CuI(pmap)]+ pmap2). In contrast,
(
(
we find the relative rates to be pmea2 ꢆ pmap2 > tmpa2,
determined qualitatively. In fact, the reduction of 10-Cl
by pmea2 in THF is too fast to be monitored by stopped-
flow kinetics even at 188 K whereas the reaction involving
tmpa2 (Figs. 13 and 14) takes greater than 5 minutes to com-
plete at 233 K. The reaction involving pmap2 at 188 K is
observed on a millisecond time-scale (Figs. 15 and 16).
The kinetic traces that we monitor for reactions involving
tmpa2 at 233 K (Fig. 14) and pmap2 at 188 K (Fig. 16) both
fit a first-order kinetic model, kobs = 1.42 · 10ꢁ5 sꢁ1 and
3 · 10ꢁ2 sꢁ1 respectively. These results are somewhat unex-
pected and for the purposes of the present study which
emphasizes qualitative aspects of the present relatively
uncommon electron-/atom- (ClÅ) transfer chemistry, fur-
ther kinetic studies have not been pursued. Future efforts
will include studies involving solvent and metal complex
The
FeIII/II
and
CuII/I
redox
waves
for
[(6L)Feꢂ ꢂ ꢂClꢂ ꢂ ꢂCu]+ (3) separate in the presence of external
substrates (i.e., CO and DCHIm), Fig. 12 and Table 2.
More specifically, addition of DCHIm to 3 results in the
CuII/I redox potential shifting negative to E1/2
=
ꢁ650 mV and a positive shift of the FeIII/II redox potential
1
to E1/2 = ꢁ490 mV occurs. Following H NMR and EPR
spectral analysis,
a
low-spin [(6L)FeII(DCHIM)2ꢂ ꢂ ꢂ
CuII(Cl)]+ species is the likely reaction product. Upon
addition of CO to this solution, the FeIII/II redox potential
Fig. 12. Cyclic voltammograms recorded in CH3CN with (Bu)4NPF6 as
the supporting electrolyte of [(6L)Feꢂ ꢂ ꢂClꢂ ꢂ ꢂCu]+ (3) (black), [(6L)FeII(D-
Fig. 13. Time-dependent UV–Vis spectra (time = 0–375 s) for the ClÅ
CHIm)2ꢂ ꢂ ꢂCuII(Cl)]
(Scheme
6d)
(gray),
and
[(6L)FeII(D-
transfer reaction of (F8)FeIII(Cl) (10-Cl) to [CuI(tmpa)(THF)]+ tmpa2) in
(
CHIm)2ꢂ ꢂ ꢂCuII(Cl)] + CO (dashed line).
THF at 233 K.

H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
1111
Fig. 14. Absorption changes as a function of time from stopped-flow UV–
Vis spectroscopy measured at 506 nm (blue), 550 nm (red), and 630 nm
(green) following the reaction of (F8)FeIII(Cl) (10-Cl) and [CuI(tm-
Fig. 16. Absorption changes as a function of time from stopped-flow UV–
Vis spectroscopy measured at 505 nm (blue), 545 nm (red), and 630 nm
(green) following the reaction of (F8)FeIII(Cl) (10-Cl) and [CuI(pmap)]+
pa)(THF)]+
(
tmpa2) in THF at 233 K. First order kinetic fits are overlaid
(
pmap2) in THF at 188 K. First order kinetic fits are overlaid as black lines,
as black lines, kobs = 1.42 · 10ꢁ5 sꢁ1
.
kobs = 3 · 10ꢁ2 sꢁ1
.
nism, the formation of a chlorine bridge between iron
and copper (Scheme 7) precedes electron-/atom- (ClÅ)
transfer. As a result, the thermodynamics calculated (vide
supra) from the free mononuclear compounds are not
necessarily relevant. Other factors are expected to play
important roles, such as those arising from bond breakage
and formation and variations in electronic coupling across
the halide bridge [47–49].
The chemical nature of this putative halide bridged
inner-sphere complex remains speculative as no intermedi-
ates were observed during the course of the reactions. The
structural changes associated with bridge formation about
copper are expected to be more significant than that at the
heme (vide supra). In this regard it is interesting to note that
the rates were fastest for pmea2 and slowest for tmpa2. As
described previously (Section 3.3), ClÅ transfer for the latter
compound involves a change in coordination number, de-
ligation of one of the pyridyl arms, and solvent coordina-
tion, Scheme 4. The latter two processes are absent for
pmea2 (and pmap2). Therefore, although it is speculative at
this time, the results are consistent with expectations that
geometric changes about copper are very important in
determining the rates of ClÅ transfer.
Fig. 15. Time-dependent UV–Vis spectra (time = 0–300 ms) for the ClÅ
transfer reaction of (F8)FeIII(Cl) (10-Cl) to [CuI(pmap)]+ pmap2) in THF at
(
188 K.
concentration variations, to elucidate further mechanistic
aspects.
3.7. Likely inner-sphere electron-/atom- (ClÅ) transfer
chemistry
While the rates for ClÅ transfer varied over impressive
ranges, from too fast for stopped-flow kinetics to millisec-
onds to minutes, by tuning the copper ligand environment,
there was no correlation with the thermodynamic driving
force estimated from cyclic voltammetry measurements.
Such correlations are well known for outer-sphere elec-
tron-transfer processes both experimentally [43,44] and
theoretically [45]. Therefore, the lack of a free energy cor-
relation suggests that there are indeed important contribu-
tions to the rates from ‘‘inner-sphere’’ pathways like those
first described by Taube [46]. For an inner-sphere mecha-
3.8. Observed rates of electron-/atom- (ClÅ) transfer and
bimolecular rate constant for CO rebinding as determined by
transient absorption spectroscopy
Due to the binucleating nature of 6L, stopped-flow UV–
Vis spectroscopy could not be utilized to measure the rates
of electron-/atom- (ClÅ) transfer. Instead, a complementary
approach, transient absorption laser flash photolysis (TA),
was employed. Here we took advantage of the well docu-
mented photodissociation of carbon monoxide from
heme–carbonyl compounds. The photorelease of CO is

1112
H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
Scheme 7.
Scheme 8.
known to occur with a quantum yield of unity in both nat-
ural and synthetic hemes [50,51].
kinetic analysis termed ‘‘decay associated spectra’’ by
Brand and coworkers [52].
We therefore prepared the compound [(6L)FeII(CO)ꢂ ꢂ ꢂ
CuII(Cl)]+ (5) by dissolution of [(6L)Feꢂ ꢂ ꢂ(Cl)ꢂ ꢂ ꢂCu]+ (3)
in CH3CN and bubbling with CO(g), Scheme 6e. Upon
excitation (kex = 532 nm; 8–10 ns FWHM) of 5 with pulsed
light in CH3CN under 1 atm CO (Scheme 8), the photoini-
tiated ejection of CO occurs and the rate at which the initial
species reforms could be measured. Typical absorption
difference spectra for such an experiment are shown at
the indicated delay times after the laser pulse in Fig. 17.
It was apparent from both the appearance wavelength
dependent bi-exponential kinetics, Eq. (1), that two differ-
ent species were photogenerated within the time response
of our apparatus (ꢄ10 ns). At wavelengths
The time resolved spectra in Fig. 17a were obtained
upon photodissociation of CO from [(6L)FeII(CO)ꢂ ꢂ ꢂ
CuII(Cl)]+ (5) and they correlate to formation of the bis-
solvento species, [(6L)FeII(CH3CN)2ꢂ ꢂ ꢂCuII(Cl)]+ (3CuCl).
Support for this assignment was gained by subtracting
the absorption spectrum of 5 from 3CuCl, Abs[(6L)-
FeII(CH3CN)2ꢂ ꢂ ꢂCuII(Cl)]+ (3CuCl) – Abs[(6L)FeII(CO)ꢂ ꢂ ꢂ
CuII(Cl)]+ (5), and overlaying the resulting spectrum on
the transient data, Fig. 17a. The excellent agreement gives
little doubt that this intermediate has been correctly
assigned. In addition, we found that the lifetime of the sol-
vent compound showed a first-order dependence on the CO
concentration. A linear plot of k1 (sꢁ1) versus [CO] yielded
a bimolecular rate constant of kCO = 1.4 · 107 Mꢁ1 sꢁ1
,
AðtÞ ¼ A1 expðꢁk1tÞ þ A2 expðꢁk2tÞ
ð1Þ
inset Fig. 17a. This value is in the same range as that mea-
sured for CO recombination to other synthetic and biolog-
ical hemes.
lower than 375 nm, the absorption transients were almost
perfectly exponential as the second state did not absorb
light appreciably at this wavelength. This allowed us to
accurately determine the value of k2. With this informa-
tion, k2 was fixed and the value of k1 was determined at
the wavelength with the highest signal-to-noise ratio,
ꢄ420 nm. With these two rate constants fixed, the wave-
length dependent absorption changes were fit to Eq. (1)
and only the amplitudes were allowed to vary. The ampli-
tudes of A1 and A2 as a function of time are shown in
Fig. 17a and 17b alone. We note that this is a standard
As a source of further comparison, the bimolecular rate
constant for CO rebinding to (F8)FeII (1) in CH3CN upon
photodissociation from (F8)FeII(CO) was also determined
here as kCO = 6.0 · 105 Mꢁ1 sꢁ1, Fig. 18. The greater than
two orders of magnitude difference between kCO for
[(6L)FeII(CO)ꢂ ꢂ ꢂCuII(ClÅ)]+ (5) versus (F8)FeII(CO) shows
that for 5, the presence of the copper(II) moiety or the dif-
6
ference in heme structure (F8 versus L) has a large effect
on the chemistry at the FeII ion in [(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+
(3CuCl) as concerns CO rebinding. One possible explanation

H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
1113
Fig. 18. Transient absorption difference spectra obtained after photo-
lysis of F8FeII(CO)(CH3CN) in CO saturated CH3CN yielding
F8FeII(CH3CN)2 (1) with various delay times monitored at 0 ls (black
squares), 10 ls (blue triangles), 25 ls (green diamonds), 50 ls (red circles),
100 ls (turquoise triangles) and 250 ls (purple stars) where the pink
calculated difference spectrum {Abs[F8FeII(CH3CN)2] (1) – Abs[F8FeII-
(CO)(CH3CN)]} is displayed for comparison in the lower inset and the
linear plot in the upper inset displays the CO dependence of this process
where the slope yields, kCO (298 K) = 6.0 · 105 Mꢁ1 sꢁ1
.
Fig. 17. Transient absorption difference spectra obtained after photolysis
of [(6L)FeII(CO)ꢂ ꢂ ꢂCuII(Cl)]+ (5) in CO saturated CH3CN yielding (a)
[(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+ (3CuCl) with various delay times monitored at 10 ns
(squares), 5 ls (circles), 10 ls (triangles), and 30 ls (diamonds) where the
red calculated difference spectrum {Abs[(6L)FeIIꢂ ꢂ ꢂCuI]+ – Abs[(6L)FeII-
(CO)ꢂ ꢂ ꢂCuII(Cl)]+ (5)} is displayed for comparison and the inset displays
the CO dependence of this process where the slope yields, kCO (298
K) = 1.4 · 107 Mꢁ1 sꢁ1, (b) [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl) with various
extrapolated delay times at 10 ns (squares), 30 ls (circles), and 60 ls
(triangles) where the blue calculated difference spectrum {Abs[(6L)-
FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl) – Abs[(6L)FeII(CO)ꢂ ꢂ ꢂCuII(Cl)]+ (5)} is displayed
for comparison.
The data for the other perhaps more interesting process
that occurs after photodissociation of CO from [(6L)FeII-
(CO)ꢂ ꢂ ꢂCuII(Cl)]+ (5) is shown in Fig. 17b. The appearance
of a new broad absorption band at ꢄ375 nm is character-
istic of an iron species in the Fe(III) formal oxidation state.
Suslick has assigned this as a ligand-to-metal charge-
transfer (LMCT) band and reported that the position of
this band can be correlated with the donor ability of the
axial ligand [30]. The experimental data observed here is
indeed in excellent agreement with the calculated difference
spectrum
of
Abs[(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+
(3FeCl
)
–
is that 6L is a better donor ligand to iron in comparison to
F8, due to the simple fact of having two less aryl fluoro-
substituents on the heme. There exists literature precedent
for this, i.e., increased rates for small molecule binding to
the heme in a heterobinuclear complex compared to bind-
ing to (F8)FeII. For example, we previously published the
bimolecular recombination of CO to [(6L)FeII(thf)2ꢂ ꢂ ꢂ
CuI(CO)]+ and (F8)FeII(thf)2 (1) following photolysis of
[(6L)FeII(CO)ꢂ ꢂ ꢂCuI(CO)]+ and (F8)FeII(CO), yielding the
second-order rate constants of 1.0 · 107 Mꢁ1 sꢁ1 and
3.2 · 106 Mꢁ1 sꢁ1 respectively [8]. Furthermore, a similar
observation concerning rates was found upon reaction of
O2 with a related Fe/Cu system, [(2L)FeIIꢂ ꢂ ꢂCuI]+, in which
Abs[(6L)FeII(CO)ꢂ ꢂ ꢂCuII(ClÅ)]+ (5), Fig. 17b, indicating
the product is 3FeCl, Scheme 8. Remarkably, the generation
of this state could not be time-resolved with our apparatus.
This indicates that chlorine atom transfer from copper to
iron occurs with a rate constant >108 sꢁ1! This rapid
rate supports the notion suggested by the stopped-flow
kinetic measurements that atom transfer occurs through
an inner-sphere mechanism with a bridging chloride
intermediate.
On a 100 ls time scale (k2), the photogenerated interme-
diate [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl) returns to a product that
contains the heme carbonyl. The spectral data were fully
consistent with transfer of the chlorine atom back to the
copper center, however the strong absorbance of the heme
did not allow us to directly monitor the coordination of the
copper center. Thus, an electron-/atom- (ClÅ) transfer pro-
2
the copper moiety of L is tridentate in nature and com-
prised of a central amine with two pyridyl-ethyl-N-donors.
In comparison to the intermolecular initial reaction of O2
with the heme only (F8)FeII (forming a heme–O2 adduct)
in a (F8)FeII/LCuI mixture, the analogous intramolecular
reaction involving the heterþobinucleating ligand (2L) to
cess takes place within an observed rate (k2) of ꢄ103 sꢁ1
.
Given that acetonitrile is a poor ligand for Cu(II), k2
(Scheme 8) is comprised of electron-/atom- (ClÅ) transfer
coupled with CO rebinding, this observed rate constant
III
I
form ½ð LÞFe ðO2ꢁÞ ꢂ ꢂ ꢂ Cu ꢃ was more than 20 times
2
faster [14].

1114
H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
does not correspond solely to single chlorine atom transfer
and cannot be directly compared to the related ClÅ transfer
the chloride bridge are likely to be important consider-
ations. Further, we suggest that significant contributions
to the observed behavior will come from copper ion struc-
tural changes occurring during the redox reactions, but dif-
fering for the various copper–ligand complexes. Ligand
effects on outer-versus inner-sphere oxidations of (ligand)-
Cu(I) complexes have been recently probed by Meyerstein
and colleagues [53]. Further studies on the present or newly
designed systems will be of interest and required to provide
additional insights.
process of the monomeric [CuI(tmpa)(CH3CN)]+ tmpa2)/
(
[(F8)FeIII(Cl)] (10-Cl) analogue. Nevertheless, we predict
that the rate of electron-/atom- (ClÅ) transfer from
[(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl) to form [(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+
(3CuCl) would exceed the corresponding value for the
(
tmpa2)/(10-Cl) reaction. Based on the trend in kobs for the
[LCuI]+/[(F8)FeII] system as obtained by stopped-flow
UV–Vis spectroscopy in relation to the [LCuII(Cl)]+ s val-
ues, copper ion moieties with square-based pyramidal
geometries, as in [(6L)FeIIꢂ ꢂ ꢂCuII(ClÅ)]+ (3CuCl), are
expected to have higher electron-/atom- (ClÅ) transfer rates,
vide supra.
Acknowledgements
We are very grateful for financial support of this re-
search (K.D.K., National Institutes of Health, GM28962
and GM60353; G.J.M., National Science Foundation,
CHE 0616500).
4. Summary/conclusions
The reactions of (F8)FeIII(Cl) (10-Cl) with a series of
[(L)CuI]+ (R2) complexes (L = TMPA, PMEA, PMAP)
lead to electron-/atom- (ClÅ) transfer chemistry, giving
(F8)FeII (1) and [(L)CuII(Cl)]+ (R20-Cl). For [CuI(tepa)]+
Appendix A. Supplementary material
CCDC 657126 contain the supplementary crystallo-
graphic data for this paper. These data can be obtained free
of charge via http://www.ccdc.cam.ac.uk/conts/retriev-
ing.html, or from the Cambridge Crystallographic Data
Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax:
(+44) 1223-336-033; or e-mail: deposit@ccdc.cam.ac.uk.
Supplementary data associated with this article can be
found, in the online version, at doi:10.1016/j.ica.2007.
08.016.
(
tepa2), a reaction with (F8)FeIII(Cl) (10-Cl) does not occur
due to unfavorable thermodynamics. Stopped-flow spec-
troscopy indicates that (F8)FeIII(Cl) (10-Cl) oxidizes the
[(L)CuI]+ (R2) complexes in THF solvent with the follow-
ing trend in relative rates [CuI(pmea)]+
(
pmea2) > [CuI(p-
tmpa2). Redox
potentials predict a differing trend, [CuI(tmpa)(CH3CN)]+
tmpa2) > [CuI(pmea)]+ pmea2) > [CuI(pmap)]+ pmap2),
map)]+
( (
pmap2) > [CuI(tmpa)(CH3CN)]+
(
(
(
where the relative driving force was estimated by calculat-
ing E1/2(oxidant) ꢁ E1/2(reductant). Thus, the results point
to a likely inner-sphere mechanism for ClÅ transfer from
heme to copper.
References
´
[1] E.E. Chufan, S.C. Puiu, K.D. Karlin, Acc. Chem. Res. 40 (2007) 563.
The study of [(6L)Feꢂ ꢂ ꢂ(Cl)ꢂ ꢂ ꢂCu]+ (3) as an analog to the
pairing of the mononuclear species (F8)FeIII(Cl) (10-Cl) and
[CuI(tmpa)(CH3CN)]+ (tmpa2) results in an interesting equi-
[2] R.A. Ghiladi, E.E. Chufan, D. delRio, E.I. Solomon, C. Krebs, B.H.
Huynh, H.w. Huang, P. Moenne-Loccoz, S. Kaderli, M. Honecker,
A.D. Zuberbuhler, L. Marzilli, R.J. Cotter, K.D. Karlin, Inorg.
Chem. 46 (2007) 3889.
[3] R.A. Ghiladi, H.W. Huang, P. Moe¨nne-Loccoz, J. Stasser, N.J.
Blackburn, A.S. Woods, R.J. Cotter, C.D. Incarvito, A.L. Rheingold,
K.D. Karlin, J. Biol. Inorg. Chem. 10 (2005) 63.
librium between [(6L)FeIII(Cl)ꢂ ꢂ ꢂCuI]+ (3FeCl
)
and
[(6L)FeIIꢂ ꢂ ꢂCuII(Cl)]+ (3CuCl). This equilibrium is displayed
as nearly 1:1 (3FeCl):(3CuCl) by 1H NMR (THF-d8) and elec-
trochemical studies (CH3CN). Upon addition of an axial
base like pyridine or 1,5-DCHIm, the equilibrium shifts
towards 3CuCl accompanied by ‘base’ binding to iron(II).
Adding carbon monoxide similarly results in an equilibrium
shift towards [(6L)FeII(CO)ꢂ ꢂ ꢂCuII(Cl)]+ (5). Upon photol-
ysis of 5, we generate 3CuCl which competitively rebinds
CO or undergoes an assumed electron-/atom- (ClÅ) transfer
process forming 3FeCl. Ultimately, 3FeCl equilibrates back to
5 in an overall reversible process. Quantifying the kinetics
was difficult due to the composite nature of the reaction.
In conclusion, electron-/atom- (ClÅ) transfer reactions
have been observed in systems involving a combination
of heme and copper ion centers, in both inter- and intra-
molecular cases. Based on the lack of correlation of reac-
tion rates and thermodynamic factors, the occurrence of
an inner-sphere chloride-bridged intermediate is suggested
although it is not observed. The details of bond-breaking
and bond-making along with electronic coupling across
´
[4] E. Kim, E.E. Chufan, K. Kamaraj, K.D. Karlin, Chem. Rev. 104
(2004) 1077.
[5] J.P. Collman, R. Boulatov, C.J. Sunderland, L. Fu, Chem. Rev. 104
(2004) 561.
[6] H.C. Fry, A.D. Cohen, J.P. Toscano, G.J. Meyer, K.D. Karlin, J.
Am. Chem. Soc. 127 (2005) 6225.
[7] R.M. Kretzer, R.A. Ghiladi, E.L. Lebeau, H.-C. Liang, K.D. Karlin,
Inorg. Chem. 42 (2003) 3016.
[8] D.W. Thompson, R.M. Kretzer, E.L. Lebeau, D.V. Scaltrito, R.A.
Ghiladi, K.-C. Lam, A.L. Rheingold, K.D. Karlin, G.J. Meyer,
Inorg, Chem. 42 (2003) 5211.
[9] A. Nanthakumar, S. Fox, K.D. Karlin, J. Chem. Soc., Chem.
Commun. (1995) 499.
[10] D.B. Rorabacher, Chem. Rev. 104 (2004) 651.
[11] M. Schatz, M. Becker, F. Thaler, F. Hampel, S. Schindler, R.R.
´
Jacobson, Z. Tyeklar, N.N. Murthy, P. Ghosh, Q. Chen, J. Zubieta,
K.D. Karlin, Inorg. Chem. 40 (2001) 2312.
[12] K.D. Karlin, J.C. Hayes, J. Shi, J.P. Hutchinson, J. Zubieta, Inorg.
Chem. 21 (1982) 4106.
[13] R.A. Ghiladi, T.D. Ju, D.-H. Lee, P. Moe¨nne-Loccoz, S. Kaderli, Y.-
M. Neuhold, A.D. Zuberbu¨hler, A.S. Woods, R.J. Cotter, K.D.
Karlin, J. Am. Chem. Soc. 121 (1999) 9885.

H.C. Fry et al. / Inorganica Chimica Acta 361 (2008) 1100–1115
1115
[14] E. Kim, M.E. Helton, S. Lu, P. Moe¨nne-Loccoz, C.D. Incarvito, A.L.
[33] I.M. Wasser, H.W. Huang, P. Moenne-Loccoz, K.D. Karlin, J. Am.
Chem. Soc. 127 (2005) 3310.
Rheingold, S. Kaderli, A.D. Zuberbuhler, K.D. Karlin, Inorg. Chem.
¨
44 (2005) 7014.
[34] E. Kim, M.E. Helton, I.M. Wasser, K.D. Karlin, S. Lu, H.-w. Huang,
P. Moe¨nne-Loccoz, C.D. Incarvito, A.L. Rheingold, M. Honecker, S.
[15] C.X. Zhang, H.-C. Liang, E.-i. Kim, J. Shearer, M.E. Helton, E. Kim,
S. Kaderli, C.D. Incarvito, A.D. Zuberbuhler, A.L. Rheingold, K.D.
¨
Kaderli, A.D. Zuberbuhler, Proc. Natl. Acad. Sci. USA 100 (2003)
¨
Karlin, J. Am Chem. Soc. 125 (2003) 634.
3623.
[16] H.C. Fry, H.R. Lucas, A.A. Narducci Sarjeant, G.J. Meyer, K.D.
Karlin, submitted for publication.
[17] R.A. Ghiladi, R.M. Kretzer, I. Guzei, A.L. Rheingold, Y.-M.
[35] B.S. Lim, R.H. Holm, Inorg. Chem. 37 (1998) 4898.
[36] D. Lancon, K.M. Kadish, J. Am. Chem. Soc. 105 (1983) 5610.
[37] G. Balducci, G. Chottard, C. Gueutin, D. Lexa, J.M. Saveant, Inorg.
Chem. 33 (1994) 1972.
Neuhold, K.R. Hatwell, A.D. Zuberbuhler, K.D. Karlin, Inorg.
¨
Chem. 40 (2001) 5754.
[38] K.M. Kadish, E. Van Caemelbecke, F. D’Souza, M. Lin, D.J. Nurco,
C.J. Medforth, T.P. Forsyth, B. Krattinger, K.M. Smith, S. Fukuz-
umi, I. Nakanishi, J.A. Shelnutt, Inorg. Chem. 38 (1999) 2188.
[39] N. Wei, N.N. Murthy, Q. Chen, J. Zubieta, K.D. Karlin, Inorg.
Chem. 33 (1994) 1953.
[18] C.A. Bignozzi, R. Argazzi, C. Chiorboli, F. Scandola, R.B. Dyer, J.R.
Schoonover, T.J. Meyer, Inorg. Chem. 33 (1994) 1652.
[19] H.C. Fry, D.V. Scaltrito, K.D. Karlin, G.J. Meyer, J. Am Chem. Soc.
125 (2003) 11866.
[20] C.X. Zhang, S. Kaderli, M. Costas, E.-i. Kim, Y.-M. Neuhold, K.D.
[40] R. Boulatov, J.P. Collman, I.M. Shiryaeva, C.J. Sunderland, J. Am.
Chem. Soc. 124 (2002) 11923.
Karlin, A.D. Zuberbuhler, Inorg. Chem. 42 (2003) 1807.
¨
[21] P. van der Sluis, A.L. Spek, Acta Cyrstallogr., Sect. A 46 (1990) 194.
[22] H.-C. Liang, E. Kim, C.D. Incarvito, A.L. Rheingold, K.D. Karlin,
Inorg. Chem. 41 (2002) 2209.
[23] W.R. Scheidt, C.A. Reed, Chem. Rev. 81 (1981) 543.
[24] F.A. Walker, U. Simonis, in: R.B. King (Ed.), Encyclopedia of
Inorganic Chemistry, 2nd ed., John Wiley & Sons Ltd., 2005, p. 2390.
[41] J.P. Collman, Y.L. Yan, J. Lei, P.H. Dinolfo, Inorg. Chem. 45 (2006)
7581.
[42] J.P. Collman, N.K. Devaraj, R.A. Decreau, Y. Yang, Y.-L. Yan,
W. Ebina, T.A. Eberspacher, C.E.D. Chidsey, Science 315 (2007)
1565.
[43] D. Rehm, A. Weller, Isr. J. Chem. 8 (1970) 259.
[44] M. Chou, C. Creutz, N. Sutin, J. Am. Chem. Soc. 99 (1977) 5615.
[45] R.A. Marcus, N. Sutin, Biochim. Biophys. Acta 811 (1985) 265.
[46] C. Creutz, P.C. Ford, T.J. Meyer, Inorg. Chem. 45 (2006) 7059.
[47] A. Haim, Prog. Inorg. Chem. 30 (1983) 273.
[48] M.Y. Udugala-Ganehenege, M.J. Heeg, L.M. Hryhorczuk, L.E.
Wenger, J.F. Endicott, Inorg. Chem. 40 (2001) 1614.
[49] J.F. Endicott, M.J. Uddin, Coord. Chem. Rev. 219 (2001) 687.
[50] E. Antonini, M. Brunori, Hemoglobin and Myoglobin in Their
Reactions, North-Holland Publishing Co., Amsterdam, 1971.
[51] T.G. Traylor, Acc. Chem. Res. 14 (1981) 102.
[52] D. Toptygin, A.M. Gronenborn, L. Brand, J. Phys. Chem. B 110
(2006) 26292.
´
[25] Z. Tyeklar, R.R. Jacobson, N. Wei, N.N. Murthy, J. Zubieta, K.D.
Karlin, J. Am. Chem. Soc. 115 (1993) 2677.
[26] R.R. Jacobson, Dissertation, State University of New York at
Albany, 1989.
[27] A.W. Addison, T.N. Rao, J. Reedijk, J. van Rijn, G.C. Verschoor, J.
Chem. Soc., Dalton Trans. (1984) 1349.
[28] N. Wei, N.N. Murthy, K.D. Karlin, Inorg. Chem. 33 (1994) 6093.
[29] B. Lucchese, K.J. Humphreys, D.-H. Lee, C.D. Incarvito, R.D.
Sommer, A.L. Rheingold, K.D. Karlin, Inorg. Chem. 43 (2004) 5987.
[30] D.N. Hendrickson, M.G. Kinnaird, K.S. Suslick, J. Am. Chem. Soc.
109 (1987) 1243.
[31] H.-C. Liang, K.D. Karlin, R. Dyson, S. Kaderli, B. Jung, A.D.
Zuberbuhler, Inorg. Chem. 39 (2000) 5884.
¨
[32] R.A. Ghiladi, K.D. Karlin, Inorg. Chem. 41 (2002) 2400.
[53] A. Burg, E. Maimon, H. Cohen, D. Meyerstein, Eur. J. Inorg. Chem.
(2007) 530.