Understanding of the mode of action of FeIII-EDDHA as iron chlorosis corrector based on its photochemical and redox behavior

Article abstract of DOI:10.1002/chem.200500286

The very low reduction potential of the chelate Fe^III-EDDHA (EDDHA= ethylenediamine N,N'-bis(2-hydroxy)phenylacetic acid) makes it unreactive in photochemically or chemically induced electron transfer processes. The lack of reactivity of this c

Full text of DOI:10.1002/chem.200500286

FULL PAPER
DOI: 10.1002/chem.200500286
Understanding of the Mode of Action of FeIII–EDDHA as Iron Chlorosis
Corrector Based on Its Photochemical and Redox Behavior**
Mar Gómez-Gallego,*[a] Daniel Pellico,[a] Pedro Ramírez-López,[a] María J. MancheÇo,[a]
Santiago Romano,[a] María C. de la Torre,[b] and Miguel A. Sierra*[a]
Abstract: The very low reduction po-
tential of the chelate FeIII–EDDHA
(EDDHA= ethylenediamine N,N’-
However, in spite of its low reduction
potential, the biological reduction of
FeIII–EDDHA is very effective. Based
on electrochemical measurements, it is
proposed that FeIII–EDDHA itself is
not the substrate of the enzyme ferric
chelate reductase. Likely, at the more
acidic pH in the vicinity of the roots,
the ferric chelate in a closed form
(FeLꢀ) could generate a vacant coordi-
nation site that leads to an open hexa-
coordinate species (FeHL) where the
reduction of the metal by the enzyme
takes place.
bis(2-hydroxy)phenylacetic
acid)
makes it unreactive in photochemically
or chemically induced electron transfer
processes. The lack of reactivity of this
complex toward light invalidates pho-
todegradation as an alternative mecha-
nism for environmental elimination.
Keywords: chelate reductase · chlo-
rosis · electrochemistry · electron
transfer · iron · photochemistry
Introduction
Polyaminecarboxylic acids have the ability to form stable
water-soluble complexes with di- and trivalent metal ions.
Accordingly, they are used to control the solubility and pre-
cipitation of metal ions in a broad range of areas, from do-
mestic products to industrial and agronomic applications.[1,2]
Owing to the increasing concern about their environmental
impact, biotransformations and mineralization processes of
the polyaminecarboxylic acids most widely used, nitrilotri-
acetic acid (NTA) and ethylenediaminetetraacetic acid
(EDTA) have been studied rather well. NTA is readily bio-
degradable,[3] but is under scrutiny due to possible adverse
health effects.[4] On the other hand, the very effective chelat-
ing agent EDTA has the disadvantage of being quite persis-
tent toward biological degradation, either as a free acid or
as a metal complex.[1] This feature has directed some atten-
tion to other mechanisms of elimination based on abiotic
processes, mainly photodegradation, oxidation by metal
oxides or hydroxides and, to a smaller degree, sorption and
sedimentation.
[a] Prof. M. Gómez-Gallego, D. Pellico, Dr. P. Ramírez-López,
Dr. M. J. MancheÇo, Dr. S. Romano, Prof. M. A. Sierra
Departamento de Química Orgµnica
Facultad de Química, Universidad Complutense
28040-Madrid (Spain)
Studies about the photoreactivity of NTA,[5–7] EDTA[8–12]
and diethylenetriaminepentaacetic acid (DTPA)[13] reveal
that photochemical degradation can be a good abiotic alter-
native pathway in the removal of these chelating agents
from water. These investigations have shown that FeIII, CuII,
and other divalent metal–NTA complexes are quite photo-
Fax : (+34)913-944-310
E-mail: margg@quim.ucm.es
sierraor@quim.ucm.es
[b] Dr. M. C. de la Torre
Instituto de Química Orgµnica, CSIC
Juan de la Cierva 3, 28006-Madrid (Spain)
[**] EDDHA= ethylenediamine N,N’-bis(2-hydroxy)phenylacetic acid.
Chem. Eur. J. 2005, 11, 5997 – 6005
ꢀ 2005 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
5997

stable, whereas FeIII–EDTA and FeIII–DTPA are photolabile
complexes, which when exposed to sunlight and in the pres-
ence of oxygen, lead to different photodegradation products
that are found to be readily biodegradable.[13,14] Other metal
complexes of EDTA (with MgII, CaII, NiII, CuII, ZnII) are un-
reactive toward sunlight.[9,10] The photoreactivity of FeIII–
EDTA and FeIII–DTPA has been explained in terms of a
light-induced reduction of the FeIII to a FeII, which is fol-
lowed by the sequential fragmentation of the ligand.
Among polyaminecarboxylic acids, those bearing pheno-
late groups are well-known products massively used as mi-
cronutrient chelating fertilizers. In particular, ethylenedia-
mine N,N’-bis(2-hydroxy)phenylacetic acid (EDDHA) is
one of the most efficient iron-chelating agents employed to
relieve iron chlorosis in plants. This nutritional disorder re-
sults in a decrease in the amount of chlorophyll in leaves
and is caused by a low absorption of iron in alkaline soil
conditions.[15] It has been estimated that in Mediterranean
countries alone, more than 4000 metric tons of FeIII–
EDDHA and other structurally related phenolic ferric che-
lating agents are used every year. In spite of this fact, their
biodegradation and photodegradation pathways are virtually
unknown and the mechanism of iron release is subject of
debate.[16]
Our current work is directed toward the synthesis of bio-
logically active bioorganometallic compounds.[17] We are in-
terested in the study of the mechanisms of reaction of
metal–organic complexes,[18] to understand the effect of both
ligand and metal in the biological action of these com-
pounds.
Herein we report a study of the photochemical behavior
of EDDHA, both as a free acid and as a FeIII complex.
Since it has been stated that the photodegradation of FeIII
polycarboxylate complexes requires the initial reduction of
the FeIII in a ligand-to-metal photoinduced electron transfer
process,[19] in this work we will also study the behavior of
the FeIII–EDDHA in electron transfer reactions, which are
either chemically or electrochemically induced. The ability
of this complex toward reduction will help us to understand
the mechanism of the iron-uptake in strategy-I plants, a pro-
cess in which the enzyme ferric chelate reductase plays a
major role.[16,20] The results of this study will allow us to
evaluate the importance of the photochemical degradation
in the removal of EDDHA and FeIII–EDDHA from the en-
vironment.[21]
Results and Discussion
The photochemical behavior of EDDHA is unknown and
for the purposes of this work is the first step towards under-
standing the photodegradation pathways and the photoprod-
ucts formed from the FeIII–EDDHA complex. EDDHA as a
free acid (isoelectric form) is insoluble in water in the pH
range of 6–7 and hence the study of its photochemical be-
havior in solution should be done either in acidic medium
(as hydrochloride 1 (see Scheme 1)) or in basic medium (as
carboxylate and phenolate salt 2 (see Scheme 2)). Solutions
of EDDHA in water at pH<2 or pH>8 are stable indefi-
nitely when stored in the dark. The irradiations with UV/Vis
light were done using a 400-W medium-pressure mercury
lamp (lmax =254, 313, 365, 436 nm). For the visible-light ex-
periments (l>313 nm) a Pyrex filter was used. The samples
were irradiated for 24 h. The irradiation of solutions of
EDDHA with UV/Vis light, in the presence of O2 and
Abstract in Spanish: El bajo valor del potencial de reducción
del quelato FeIII-EDDHAle convierte en un compuesto
prµcticamente inerte en procesos de transferencia electrónica,
tanto inducida por vía química como por vía fotoquímica.
Este hecho elimina la fotodegradación como una posible al-
ternativa abiótica para la destrucción del FeIII-EDDHAen el
medio ambiente. Sin embargo, a pesar de su bajo potencial
de reducción, el FeIII-EDDHAes uno de los mejores correc-
tores de clorosis fØrrica que se conocen, lo que implica que el
proceso de reducción biológica de este complejo se produce
de forma eficaz. En este trabajo, basµndonos en medidas
electroquímicas, proponemos que el complejo FeIII-EDDHA
en su forma (FeLꢀ) no es el sustrato de la enzima quelato re-
ductasa fØrrica. Posiblemente, debido a la acidificación que
se produce en la proximidad de las raíces, el complejo FeLꢀ
podría generar una vacante de coordinación y adoptar una
forma hexacoordinada abierta (FeHL) en la que se produce
la reducción enzimµtica.
Scheme 1.
5998
www.chemeurj.org
ꢀ 2005 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Chem. Eur. J. 2005, 11, 5997 – 6005

Photochemistry of FeIII–EDDHA
FULL PAPER
under acidic conditions (pH 1), leads to o-hydroxyphenyl
glycine (3) as the main reaction product (44%), together
with unaltered starting material (51%) (Scheme 1). The
photodegradation is much slower in visible light (l>
313 nm) and only trace amounts of 3 were obtained after ir-
radiation for 24 h (93% of EDDHA was recovered unal-
tered in the experiment). At basic pH values (pH 12), the
main isolated photoproduct on both UV/Vis and visible irra-
diation is salicylaldehyde imine 4 (11%), accompanied by
73–75% of recovered unreacted starting material
(Scheme 2). Additionally, hydrolysis of the imine 4 is ob-
served, and salicylaldehyde together with its oxidation prod-
uct (salicylic acid) were also identified in the crude reaction
products. The same product distribution but in considerably
lower yields was observed when the experiments were car-
ried out in deaerated solutions. The irradiation of suspen-
sions of EDDHA in water at pH 7 afforded the unaltered
product after reaction times of 24 h under all conditions es-
sayed.
24 h. Trace amounts of 3 (< 2%) were detected in the reac-
tion product. In basic medium, only small amounts (6%) of
salicylic acid and salicylaldehyde were identified as degrada-
tion products. Finally, in neutral solution, the product was
totally inert to photodegradation.
The results obtained in the different experiments show
that FeIII–EDDHA is very resistant to photochemical degra-
dation. Solutions of the complex in water are almost totally
inert to light irradiation, even when the experiment was car-
ried out at pH values well above or below the pH range in
which the iron chelate is stable. In those cases, the small
amount of photoproducts detected should come from the
photodegradation of the free ligand EDDHA, which is pres-
ent in solution in equilibrium with the complex.[24]
The behavior of FeIII–EDDHA toward light contrasts
with that previously reported for the very photoreactive
FeIII–EDTA and FeIII–DTPA.[7,12,13] In those cases, a mecha-
nism analogous to that proposed for the photochemistry of
FeIII–oxalate[25] and FeIII–citrate complexes[26] has been sug-
gested (Scheme 3). After the excitation of the complex with
These results indicate that the chelating agent EDDHA is
rather stable to photodegradation. Only at extreme pH
values, well above or below the typical environmental pH
range (4 to 8), is some photoreactivity observed. The pro-
cess is clearly pH-dependent and requires UV light. Under
ꢀ
strong acidic conditions, 3 is formed, presumably by a-C H
fragmentation of the amino bonds followed by deprotona-
tion and disproportionation, to finally form the unstable
glyoxal imine 5, which is hydrolyzed in the acidic medium
(Scheme 1).[22] This type of reactivity is similar to that re-
ported for DTPA, which after successive fragmentations
leads to glycine as one of the main photodegradation prod-
ucts.[13] In a strong basic medium, EDDHA is less reactive.
Salicylaldehyde imine 4 is formed in low yield by photoin-
duced decarboxylation of 2, a well-known process for a-
amino carboxylates (Scheme 2).[23]
Scheme 3.
light, the FeIII is reduced by electron transfer (SET) from
one of the carboxylate groups of the ligand, leading to FeII
and a carboxylate radical cation species from which the pho-
tofragmentation products could be formed. The irreversible
oxidation of ferrous iron in the medium finally leads to the
precipitation of iron hydroxide.
Our experimental results show that the photolytic ligand-
to-metal electron transfer does not occur in FeIII–EDDHA,
at least not in a productive way. Maybe the ligand already
transfers an electron to the metal (SET) and the radical
cation is formed, but a fast back electron transfer reaction
(BET) re-forms FeIII, which impeding the progress of the re-
action.[27] On the other hand, the formation of light-induced
carboxylate radicals has been proposed for oxalates, malo-
nates, and citrates, but even though this a very likely process
for phenol-polyaminecarboxylic acids, there are no data re-
ported about this reaction for such compounds.
It has been stated that there is a strong relationship be-
tween the photosensitivity of a metal chelate and the struc-
ture of the ligand.[28] Generally, transition-metal chelates
have intense absorption bands assigned to transitions involv-
ing charge-transfer excitations. However, chelate ring chro-
mophores such as iron phenanthroline and other related
complexes are essentially insensitive to light despite their in-
tense UV/Vis absorptions. This fact indicates that light ab-
sorption is not the only factor that determines the photo-
reactivity of a metal complex.
Scheme 2.
The chelate FeIII–EDDHA is stable in the pH range 3–
10.[24] The photochemical behavior of this compound has
been studied at pH 2, 6.5, and 11, in the presence of O2 and
with UV or visible light. In acidic solution, the iron chelate
was recovered practically unaltered after irradiation for
As we have discussed above, ligand-to-metal photoin-
duced electron transfer does not take place in the FeIII–
Chem. Eur. J. 2005, 11, 5997 – 6005
ꢀ 2005 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
www.chemeurj.org
5999

M. Gómez-Gallego, M. A. Sierra et al.
Table 1. Electrochemical data obtained for o,o-EDDHA.
EDDHA complex. However, the mechanism of iron release
to plants from this type of chelate requires the reduction of
FeIII to FeII. Accordingly, we decided to explore the behav-
ior of the FeIII–EDDHA complex in the presence of chemi-
cal electron transfer reagents such as Na or Li naphthale-
nide,[29] samarium iodide (SmI2),[30] or potassium graphite
(C8K).[31] The reactions were carried either in solution
(naphthalenides, SmI2) or in heterogeneous phase (C8K). In
all cases, the complex was recovered unaltered after the pro-
cess. This fact was confirmed by ESI-MS of the reaction
product, which showed a peak at m/z 412 [MꢀH]ꢀ (nega-
tive-ion mode) corresponding to unaltered FeIII–EDDHA. If
the reduction of the metal in the reaction medium had oc-
curred, free FeII should have been released to the solution.
Considering that the affinity of the EDDHA for FeII is very
low (logK=14.30),[32] the free ligand should have been ob-
served after the reaction, which was not the case. The reac-
tivity of free EDDHA in the presence of the electron trans-
fer reagents under the conditions essayed for the complex
was also checked, but the product was recovered unaltered
in all cases.
pH
Epa [V]
pH
7.4
8.1
9.3
10.1
11.1
Epa [V]
2.0
3.0
4.0
5.0
6.6
1.12
1.07
0.98
0.93
0.80
0.73
0.78
0.75
0.67
0.64
values are summarized in Table 1. In all cases, they show a
clear oxidation wave attributable to the oxidation of the
phenol group. The higher Epa values correspond to the more
acid pH values, but as the pH of the medium increases, the
phenol groups are deprotonated and the Epa values consis-
tently decrease.
The reduction potential of the FeIII–EDDHA/FeII–
EDDHA couple can be estimated by applying the Nernst
[33–35]
equation to the system [Eq. (1)],
where bIII and bII are
the stability constants of the FeIII–EDDHA and FeII-
EDDHA complexes, respectively, and E8 is 0.732 V.
EoꢀE ¼ 0:0591 log ðbIII=bIIÞ
ð1Þ
Considering logbIII 35.09[23] and logbII 14.30,[32] the calcu-
lated value of the reduction potential for the FeIII–EDDHA/
FeII–EDDHA couple is E=ꢀ0.497 V, which is significantly
lower than those reported for the FeIII/FeII–EDTA (E=
+0.120 V) and FeIII/FeII–DTPA (E=+0.030 V) couples.[33]
Therefore, it is not surprising that FeIII–EDDHA, FeIII–
EDTA, and FeIII–DTPA showed very different behavior
toward chemical and photochemical reduction. The conse-
quences on the mode of action of those compounds should
differ accordingly.
The cyclic voltammogram for FeIII–EDDHA was regis-
tered and is depicted in Figure 2. It shows quasi-reversible
electrochemistry with a half wave potential E1/2 =ꢀ 0.560 V,
[E1/2 =(Epa + Epc)/2], even lower than the calculated value
(ꢀ0.497 V). The very negative reduction potential reveals
the greatly higher selectivity of EDDHA for FeIII over FeII,
Redox behavior of FeIII–EDDHA: The nature of the ligand
in iron chelates is a key factor in the redox behavior of the
complex. In fact, it is known that the extreme variability of
the FeIII/FeII redox potential can be finely tuned by well-
chosen ligands.[33] The standard redox potential of the FeIII/
FeII couple in water is +0.732 V.[34] However, the range of
redox potentials of the LFeIII/LFeII couples can vary from
ꢀ1.0 V to +1.0 V.[33] Thus, iron can encompass almost the
entire biologically significant range of redox potentials, from
ꢀ0.40 V for semiflavin/dihydroflavin to +0.46 V for the
ꢀ
Fenton reagent or to +0.94 V for the O2 /H2O2 system. As
far as we are aware, the redox properties of FeIII–EDDHA
have not been reported.
Redox properties of the free ligand EDDHA were studied
first in order to evaluate its ability as electron donor in an
electron transfer process. The cyclic voltammograms of
1 mm solutions of EDDHA at pH values ranging from 5.0 to
9.3 are displayed in Figure 1 and the corresponding Epa
Figure 1. Cyclic voltammograms of the free ligand EDDHA (1 mm) in
0.1m buffered phosphate solutions.
Figure 2. Cyclic
voltammogram
of
FeIII–EDDHA.
Conditions:
ꢀ3
[FeCl3·6H2O ] 10 m, [o,o-EDDHA] 10ꢀ2 m, pH 7.4.
6000
www.chemeurj.org
ꢀ 2005 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Chem. Eur. J. 2005, 11, 5997 – 6005

Photochemistry of FeIII–EDDHA
FULL PAPER
as reflected in the stability constants of the corresponding
complexes. If reduction potentials are a good measurement
of the electron-accepting ability of FeIII complexes, the very
low reduction potential of FeIII–EDDHA explains its experi-
mentally proved reluctance toward photoreduction and also
to chemically induced reduction processes.
All these experimental data are of great importance to
understand the mechanism of iron release from FeIII–
EDDHA in soils. Iron uptake by strategy-I plants involves
the reduction of FeIII to FeII by the NADPH-dependent
enzyme ferric chelate reductase.[16,20] However, when consid-
ering the experimental value of the redox potential of this
complex (E=ꢀ0.560 V), the reduction of FeIII–EDDHA by
a
NADPH-enzyme
(E(NADPH/NADP+)=ꢀ0.324 V)
should not take place. This is very odd, as FeIII–EDDHA is
one of the most efficient correctors of iron chlorosis
known.[23,36] It can be argued that the measured values of
redox potentials should be considered cautiously, as they
have been obtained in experimental conditions that are not
fulfilled within cells. Although this is true, it is also reasona-
ble to assume that the differences between calculated and
real values should be small.[33]
Scheme 4.
this process could be induced by the lowering of the pH in
the roots by the proton-pumping ATPase.
To support this hypothesis we have studied the redox be-
havior of the complex FeIII–EDDHA at different acid pH
values ranging from 3.1 to 6.0. The cyclic voltammograms
are displayed in Figure 3. At the most acid pH values they
A similar oddity in the reduction by biological reducing
agents has been remarked for some ferrisiderophores, the
chelating molecules released by bacteria and fungi growing
under iron stress.[37] In these cases, it has been proposed that
the thermodynamically unfavorable FeIII to FeII biological re-
duction can be understood if a chelating agent Y, with
strong affinity for FeII, was already present in the medium.
Ferrisiderophores are released to bind FeIII and the entire
FeIII–siderophore complexes are taken up back into the cell,
where the reduction occurs. Then, complexing agent Y is
able to take FeII from the resulting FeII–siderophore com-
plex, since siderophores have little affinity for FeII. The bio-
logical reduction is driven to completion as now depends on
the FeIII–siderophore/FeII Y quotient (see [Eq. (1)]). This
model cannot be applied to strategy I plants in which the
iron is taken from extracellular ferric chelates and the re-
duction FeIII to FeII by the enzyme takes place outside the
cell.
Taking these points into consideration, the question is
how the ferric chelate reductase is able to reduce FeIII to
FeII in FeIII–EDDHA. It is known that strategy I plants in-
crease their capacity for taking up iron by excreting protons
from the root surface by activation of an ATPase.[16,20] Thus,
it could be postulated that at the acid pH of the rhizosphere
(pH 5), the enzymatic reduction does not take place at the
FeIII–EDDHA chelate itself, but at a different species were
the reduction of the metal occurs more easily. A recent
study has shown that in a wide pH range (4–10) the predom-
inant form of FeIII–EDDHA in solution is the closed octahe-
dral species FeLꢀ (Scheme 4). However, at acid pH values it
could be also in the form of the hexacoordinate open spe-
cies (FeHL) (Scheme 4).[24] The formation of FeHL requires
the generation of a vacant coordination site that is replaced
by a water molecule, leaving one of the strong binding
points of the chelate (the phenolic oxygen atom) free. All
Figure 3. Cyclic voltammograms of FeIII–EDDHA. Conditions:
ꢀ3
[FeCl3·6H2O ] 10 m, [o,o-EDDHA] 10ꢀ2 m in 0.1m buffered phosphate
solutions. a) pH 3.1 and 4.1; b) pH 5.0 and 6.0.
Chem. Eur. J. 2005, 11, 5997 – 6005
ꢀ 2005 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
www.chemeurj.org
6001

M. Gómez-Gallego, M. A. Sierra et al.
have a clear reduction wave at Epc ꢀ0.30 V (pH 3.1) and Epc
ꢀ0.51 V (pH 4.1), whereas at higher pH values they show
quasi-reversible reduction waves at E1/2 ꢀ0.37 V (pH 5.0)
and E1/2 ꢀ0.48 V (pH 6.0). The electrochemical data are
summarized in Table 2. The experimental results clearly
Table 2. Reduction potentials of FeIII–EDDHA complexes at the differ-
ent pH essayed.
Compound
pH
Epc [V]
E1/2 [V]
FeIII–o,o-EDDHA
FeIII–o,o-EDDHA
FeIII–o,o-EDDHA
FeIII–o,o-EDDHA
FeIII–o,o-EDDHA
FeIII–o,p-EDDHA
3.1
4.1
5.0
6.0
7.4
5.1
ꢀ0.30
ꢀ0.51
ꢀ0.54
ꢀ0.57
ꢀ0.61
ꢀ0.30
ꢀ0.37
ꢀ0.48
ꢀ0.56
ꢀ0.22
Figure 4. Cyclic voltammogram of FeIII–o,p-EDDHA. Conditions:
ꢀ3
[FeCl3·6H2O ] 10 m, [o,p-EDDHA] 10ꢀ2 m in 0.1m buffered phosphate
solution, at pH 5.
demonstrate that the ability of FeIII–EDDHA toward reduc-
tion increases with the lowering of the pH in the medium.
The higher concentration of the open species FeHL in the
solution at lower pH values should be responsible for the
systematic increase of the reduction potentials observed in
Table 2.
The assumption that an open species such as FeHL could
be the key to understanding the enzymatic reduction of the
ferric chelate is also supported by a recent study about the
structure and stability of the complex Fe–o,p-EDDHA, a
positional isomer of FeIII–EDDHA.[38] In this octahedral
complex, the p-hydroxy phenolate group is unable to bind
FeIII, and depending on the pH, a water molecule (up to
pH 6.3) or a hydroxy group (from pH 6.3 to 9.2) occupies
the vacant coordination position (Scheme 4). In spite of this
fact, it has been demonstrated that this complex is as effi-
cient as Fe–o,o-EDDHA as iron chlorosis corrector. The
structure of Fe–o,p-EDDHA (Scheme 4) is very similar to
that of the open form FeHL and thence this complex is an
excellent model to study the electrochemical behavior of the
FeHL species. The cyclic voltammogram of Fe–o,p-EDDHA
at pH 5.0 (X=H2Oin Scheme 4) is displayed in Figure 4
and the data are summarized in Table 2. It shows a quasi re-
versible reduction wave at E1/2ꢀ0.22 V, a value considerably
higher than that measured for the FeIII–EDDHA in the
FeLꢀ form (E1/2 ꢀ0.56 V). The comparison between these
two values provides an indication about the very different
ability of the species FeHL (open) and FeLꢀ (closed octahe-
dral) toward reduction.
Accordingly, we can propose a model that explains the re-
duction of FeIII–EDDHA by the enzyme ferric chelate re-
ductase. At the more acid pH in the vicinity of the roots, the
complex FeIII–EDDHA in the open form FeHL is reduced
by the enzyme. Free FeII and EDDHA are released in the
process and finally, the ferrous ion may be transported into
the cell by a strong chelator Y (present in the cell mem-
brane) in the form of a FeIIY complex (Scheme 5). The com-
plex FeIII–o,p-EDDHA is already in the required FeHL
open form and can be reduced directly by the enzyme.
Scheme 5.
Conclusion
In this work, we have demonstrated that FeIII–EDDHA is
very persistent toward photodegradation. This is due to the
very low reduction potential of the complex, which makes it
unreactive in photochemically or chemically induced elec-
tron transfer processes. In spite of these facts, the biological
reduction of FeIII–EDDHA by a ferric chelate reductase is
very effective. The reduction does not take place on the
complex in the FeLꢀ octahedral closed form, but on a hexa-
coordinate open species (FeHL) formed at the acid pH of
the rizhosphere. The formation of this species requires the
generation of a vacant coordination site in the FeLꢀ com-
plex that is filled with a water molecule. Our electrochemi-
cal results demonstrate that the reduction of the complex in
the open form FeHL should be easier than in the octahedral
closed form FeLꢀ. Further investigation on the photochem-
istry and redox properties of other EDDHA-like iron com-
plexes structurally related to FeIII–o,p-EDDHA are current-
ly in progress in our laboratories.
6002
www.chemeurj.org
ꢀ 2005 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Chem. Eur. J. 2005, 11, 5997 – 6005

Photochemistry of FeIII–EDDHA
FULL PAPER
aqueous fraction. The organic phase yielded o-hydroxyphenylglycine 3
(15 mg; 15%).
Experimental Section
Irradiation at pH 12: A degassed solution of 2 (100 mg; 0.27 mmol) in
distilled water (45 mL) was irradiated for 24 h. After a workup analogous
to that described above, imine 4 (10 mg; 13%) was obtained from the or-
ganic phase. The aqueous phase yielded pure o,o-EDDHA (83 mg;
83%).
General: Analytically pure o,o-EDDHA was prepared following our pre-
viously reported synthetic method.[24,39] The complex FeIII–EDDHA was
obtained as a sodium salt from the pure ligand o,o-EDDHA and ferric
chloride in basic medium following the previously reported procedure.[40]
Analytically pure o,p-EDDHA was prepared following our previously re-
ported method.[38b] Reactions requiring an inert atmosphere were con-
ducted under argon and the glassware was flame-dried under vacuum
(<0.2 mmHg). Tetrahydrofuran was distilled from sodium and benzophe-
none immediately prior to use. 1H NMR and 13C NMR spectra were re-
corded at 228C on Bruker Avance 300 (300.1 and 75.4 MHz) or Bruker
200-AC (200.1 and 50 MHz) spectrometers. Chemical shifts are given in
ppm relative to TMS (1H, 0.0 ppm), CDCl3 (13C, 77.0 ppm), C3D6O (1H,
2.0 ppm), C3D6O (13C, 206.0 ppm), D2O/NaCO3 (1H, 4.78 ppm), D2O/
NaCO3 (13C, 165.7 ppm). ESI-MS spectra were carried out in methanol
using an ESQUIRE-LC (Bruker Daltonic, Bremen, Germany) ion trap
spectrometer by using the negative-ion mode. The stainless-steel capillary
was held at a potential of 5.0 kV. Nitrogen was used as nebulizer gas at a
flow-rate of 3.98 Lminꢀ1 (nebulizer pressure 11 psi) at 1508C.
Visible irradiations (l > 313 nm) of EDDHA in the presence of O2
Irradiation at pH 1: A solution of 1 (100 mg; 0.27 mmol) in distilled
water (45 mL) was irradiated for 24 h. After a workup analogous to that
described above, pure o,o-EDDHA (93 mg; 93%) was recovered unal-
tered from the aqueous phase. Trace amounts of o-hydroxyphenylglycine
3 were detected by 1H NMR spectroscopy in the residue obtained from
the organic phase.
Irradiation at pH 12: A solution of 2 (400 mg; 1.1 mmol) in distilled
water (45 mL) was irradiated for 24 h. After a workup analogous to that
described above, imine 4 (45 mg; 11%) was obtained from the organic
phase and pure o,o-EDDHA (300 mg; 75%) was recovered unaltered
from the aqueous phase.
Irradiation at pH 7: A suspension of EDDHA (isoelectric form) (400 mg;
1.1 mmol) in distilled water (45 mL) was irradiated for 48 h. The suspen-
sion was filtered to yield unreacted EDDHA (280 mg; 70%). The aque-
ous layer was extracted with CH2Cl2, the extracts were dried (MgSO4),
filtered, and evaporated to dryness at reduced pressure, to yield trace
Photochemical procedures: All photochemical reactions were carried out
in distilled water at 228C in a magnetically stirred Pyrex vessel using a
400-W medium-pressure mercury lamp (lmax 254, 313, 365, 436 nm)
placed in a water-cooled quartz immersion well. For the visible-light ex-
periments (lmax 313 nm) a Pyrex filter was used. The pH of the different
experiments was adjusted with HCl 2m or NaOH 3m solutions in distilled
water. The samples were irradiated either in the presence or in the ab-
sence of O2. In the experiments with O2, air was bubbled through the sol-
utions during the irradiation. In the deoxygenated experiments the sam-
ples were purged with argon for at least 20 min prior to the irradiation
and then irradiated in sealed tubes.
1
amounts (< 2%) of salicylaldehyde (determined by H NMR spectrosco-
py). The aqueous layer was evaporated at reduced pressure to yield un-
reacted o,o-EDDHA (106 mg; 26%) ) (overall yield 96%).
Visible irradiations (l>313 nm) of EDDHA in the absence of O2
Irradiation at pH 1: A degassed solution of 1 (100 mg; 0.27 mmol) in dis-
tilled water (45 mL) was irradiated for 2 h. After a workup analogous to
that described above, o,o-EDDHA was quantitatively recovered from the
aqueous fraction No products were obtained from the organic phase.
UV/Vis irradiations of EDDHA in the presence of O2
Irradiation at pH 1: A solution of 1 (100 mg, 0.27 mmol), in distilled
water (45 mL) was irradiated for 24 h. The crude mixture was extracted
at pH 5 with CH2Cl2, the extracts were dried (MgSO4), filtered, and
evaporated to dryness at reduced pressure, to yield o-hydroxyphenylgly-
cine 3 (44 mg; 44%): solid; 1H NMR (200 MHz, C3D6O): d=7.34–7.31
(d, J=7.5 Hz, 1H; ArH), 7.20–7.11 (t, J=8.5 Hz, 1H; ArH), 6.88–6.81 (t,
J=7,5 Hz, 2H; ArH), 5.46 ppm (s, 1H); 13C NMR (C3D6O): d=70.1,
116.5, 120.2, 129.1, 129.8, 155.8, 174 ppm. The aqueous layer was evapo-
rated at reduced pressure to yield unreacted o,o-EDDHA (50 mg; 51%):
beige solid; 1H NMR (300 MHz, D2O/NaCO3): d=7.26–7.13 (m, 4H;
ArH), 6.88–6.78 (m, 4H; ArH), 4.44 (s, 1H), 4.37 (s, 1H), 2.94–2.75 ppm
(m, 4H); 13C NMR (300 MHz, D2O/NaCO3): d=44.4, 63.3, 116.9, 117.5,
123.5, 128.7, 128.8, 158.1, 177.3 ppm.
Irradiation at pH 12: A degassed solution of 2 (100 mg; 0.27 mmol) in
distilled water (45 mL) was irradiated for 24 h. After a workup analogous
to that described above, o,o-EDDHA was quantitatively recovered from
the aqueous fraction No products were obtained from the organic phase.
UV/Vis irradiations of FeIII-EDDHA in the presence of O2
Irradiation at pH 2: A solution of the sodium salt of FeIII–EDDHA
(900 mg) in distilled water (200 mL) was irradiated for 24 h. The solution
was filtered to remove any iron oxides formed and extracted with
CH2Cl2. The extracts were dried (MgSO4), filtered, and evaporated to
dryness at reduced pressure to yield o-hydroxyphenylglycine (3) (15 mg;
2%) (identified by 1H NMR spectroscopy). The aqueous layer was
evaporated at reduced pressure and a purple solid [Na(Fe–EDDHA)]
(702 mg; 78%) was obtained. The analysis of the product was made by
iron extraction following a reported procedure:[41] The chelate was dis-
solved in a deaerated 3m KOH solution (15 mL) and was left to stand at
room temperature for 20 min in the dark. The precipitated Fe(OH)3 was
removed by centrifugation, the clean solution was acidified at pH 6, and
the solvent was evaporated under reduced pressure. The resulting beige
solid was analyzed by NMR spectroscopy and identified as o,o-EDDHA
by comparison of the signals of the NMR spectra with those of an au-
thentic sample.
Irradiation at pH 12: A solution of 2 (100 mg, 0.27 mmol) in distilled
water (45 mL) was irradiated for 24 h. After a workup analogous to that
described above, imine 4 (11 mg; 11%) and unreacted o,o-EDDHA
(73 mg; 73%) were obtained.
Imine 4: dark yellow solid; 1H NMR (200 MHz, CDCl3): d=8.26 (s, 1H;
CH=N), 7.27–7.14 (m, 2H; ArH), 6.90–6.75 (m, 2H; ArH), 3.86 ppm (s,
2H); 13C NMR (200 MHz, CDCl3): d=59.6, 116.8, 118.5, 131.3, 132.2,
160.9, 166.4 ppm.
Irradiation at pH 11: A solution of the sodium salt of FeIII–EDDHA
(900 mg) in distilled water (200 mL) was irradiated for 24 h. A workup
analogous to that described above yielded unaltered Na(Fe–EDDHA)
(650 mg; 72%) together with a 1:6 mixture of salicylaldehyde and salicyl-
Irradiation at pH 7: A suspension of EDDHA (isoelectric form) (400 mg;
1.1 mmol) in distilled water (45 mL) was irradiated for 48 h. The suspen-
sion was filtered to yield unreacted EDDHA (250 mg; 62%). The aque-
ous layer was extracted with CH2Cl2, the extracts were dried (MgSO4),
filtered, and evaporated to dryness at reduced pressure, to yield a 1:2
mixture of imine 4 and salicylaldehyde (10 mg; ratio determined by
1H NMR spectroscopy). The aqueous layer was evaporated at reduced
pressure to yield unreacted o,o-EDDHA (140 mg; 35%) (overall yield
97%).
1
ic acid (63 mg; 7%) (analyzed by H NMR spectroscopy).
Irradiation at pH 6.5: A solution of the sodium salt of FeIII–EDDHA
(1.0 g) in distilled water (200 mL) was irradiated for 72 h. A workup anal-
ogous to that described above yielded Na(Fe–EDDHA) in quantitative
yield.
Visible irradiations (l > 313 nm) of FeIII–EDDHA in the presence of O2
UV/vis
irradiations
of
EDDHA
in
absence
of
O2
Irradiation at pH 2: A solution of the sodium salt of FeIII–EDDHA
(900 mg) in distilled water (200 mL) was irradiated for 24 h. A workup
analogous to that described above yielded Na(Fe–EDDHA) unaltered
Irradiation at pH 1: A solution of 1 (100 mg; 0.27 mmol) in distilled
water (45 mL) was irradiated for 24 h. After a workup analogous to that
described above, o,o-EDDHA (80 mg, 80%) was recovered from the
Chem. Eur. J. 2005, 11, 5997 – 6005
ꢀ 2005 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
www.chemeurj.org
6003

M. Gómez-Gallego, M. A. Sierra et al.
(823 mg; 91%) and a 1:2:2 mixture of o-hydroxyphenylglycine (3), salicy-
laldehyde, and salicylic acid (18 mg; 2%) (analyzed by 1H NMR spectros-
copy).
periments with Na(Fe-EDDHA), analysis of the product by ESI mass
spectrometry revealed only the peak m/z 412 [MꢀH]ꢀ (negative-ion
mode) for the unaltered product.
Irradiation at pH 11: A solution of the sodium salt of FeIII–EDDHA
(900 mg) in distilled water (200 mL) was irradiated for 24 h. A workup
analogous to that described above yielded Na(Fe–EDDHA) unaltered
(700 mg; 78%) together with a mixture of salicylaldehyde and salicylic
The same results were obtained when lower excesses (3:1) of C8K/Fe-
EDDHA were employed.
Electrochemical measurements: Cyclic voltammetric experiments were
performed at room temperature in 0.1m phosphate buffer. A Metrohm
6.084.010 glassy carbon electrode (GCE) was used as working electrode.
A BAS MF 2063 Ag/AgCl 3m reference and a Pt wire counter electrode
were employed. All voltammetric measurements were carried out using a
PGSTAT 12 potentiostat from Autolab. The electrochemical software
was the General Purpose Electrochemical System (GPES) (EcoChemie
B.V.).
1
acid (56 mg; 6%) (analyzed by H NMR spectroscopy).
Irradiation at pH 6.5: A solution of the sodium salt of FeIII–EDDHA
(1.0 g) in distilled water (200 mL) was irradiated for 24 h. A workup anal-
ogous to that described above yielded quantitatively the starting Na(Fe–
EDDHA) (trace amounts of imine 4 were detected).
Electron transfer reactions
General procedure for the reactions with Na/Li naphthalenide:[29] In a
two-neck round-bottom flask equipped with reflux condenser and mag-
netic stirring bar, sodium and naphthalene in dry THF were stirred under
argon for 4 h at room temperature (during which time the solution
became dark green). The solution was transferred by using a cannula to a
flask containing the chelate Na(Fe–EDDHA) or the free ligand (o,o-
EDDHA) and stirred at room temperature overnight. The ratio chelate
(or ligand)/sodium/naphthalene was 1:7.5:7.8. The reaction was quenched
with water at 08C and then filtered. The solid was dried and analyzed. In
the experiments with the free ligand o,o-EDDHA, this product was re-
covered unaltered as shown by NMR spectroscopy. In the experiments
with Na(Fe–EDDHA), the product was analyzed by ESI mass spectrom-
etry, which revealed only the peak m/z 412 [MꢀH]ꢀ (negative-ion mode)
for the unaltered product.
General procedure for the reactions of with SmI2:[30] A solution of CH2I2
in anhydrous THF was added dropwise at room temperature to a stirred
suspension of Sm in anhydrous THF under argon. The mixture was stir-
red until the color turned dark green (3 h) after which time most of the
SmIII had been consumed. The reaction mixture was stirred for an addi-
tional 30 min and transferred by using a cannula to a flask containing the
chelate Na(Fe–EDDHA) or the free ligand (o,o-EDDHA). The ratio
complex (or ligand)/Sm/CH2I2 was 1:2:1. The mixture was allowed to stir
at room temperature overnight. The reaction was quenched and the fil-
trate was washed with CH2Cl2, dried, and analyzed.
Acknowledgements
Financial support by the Spanish Ministerio de Ciencia y Tecnología
(Grants No. BQU2001–1283 and CTQ-2004–0650) and Comunidad de
Madrid (Grant No. 07M-0043–2002) are gratefully acknowledged.
[1] M. Bucheli-Witschel, T. Egli, FEMS Microbiol. Rev. 2001, 25, 69–
106.
[2] S. Tandy, K. Bossart, R. Mueller, J. Ritschel, L. Hauser, R. Schulin,
B. Nowack, Environ. Sci. Technol. 2004, 38, 937–944.
[3] R. L. Anderson, E. B. Bishop, R. L. Campbell, Crit. Rev. Toxicol.
1985, 15, 1–102.
[4] Y. Ebina, S. Okada, S. Hamazaki, F. Ogino, J. L. Li, O. Midorikawa,
J. Natl. Cancer Inst. 1986, 76, 107–113.
[5] R. J. Stolzberg, D. N. Hume, Environ. Sci. Technol. 1975, 9, 654–
656.
[6] C. H. Langford, M. Wingham, V. S. Sastri, Environ. Sci. Technol.
1973, 7, 820–822.
[7] A. Svenson, L. Kaj, H. Bjçrndal, Chemosphere 1989, 18, 1805–1808.
[8] S. S. Jones, F. A. Long, J. Phys. Chem. 1952, 56, 25–33.
[9] H. B. Lockhart Jr., R. V. Blakeley, Environ. Sci. Technol. 1975, 9,
1035–1038.
[10] P. Natarajan, J. F. Endicott, J. Phys. Chem. 1973, 77, 2049–2054.
[11] G. Karametaxas, S. J. Hug, B. Sulzberger, Environ. Sci. Technol.
1995, 29, 2992–3000.
[12] F. G. Kari, S. Hilger, S. Canonica, Environ. Sci. Technol. 1995, 29,
1008–1017.
Reactions of o,o-EDDHA with SmI2: The formation of SmI2 was accom-
plished as described above. The reaction was carried out either in THF,
THF/MeOH (2:1), or THF/H2O(1:1) mixtures using o,o-EDDHA/SmI2
ratios of 1:2.2, 1:2.9, and 1:5. The reaction was quenched by addition of a
few drops of 0.1m HCl or phosphate buffer (pH 8). The analysis by
NMR spectroscopy of the filtrate indicated that o,o-EDDHA was recov-
ered unaltered in all cases.
[13] S. Metsarinne, P. Rantanen, R. Aksela, T. Tuhkanen, Chemosphere,
2004, 55, 379–388.
[14] B. Nowack, U. Baumann, Acta Hydrochim. Hydrobiol. 1998, 26,
104–108.
Reactions of Na[Fe-EDDHA] with SmI2: The formation of SmI2 was ac-
complished as described above. The reaction was carried out either in
THF, THF/MeOH (2:1), or THF/H2O(1:1) mixtures using Na(Fe-
EDDHA)/SmI2 ratios of 1:2.2 and 1:5. The reaction was quenched by ad-
dition of a few drops of 0.1m HCl or phosphate buffer (pH 8). Analysis
of the product by ESI mass spectrometry revealed only the peak at m/z
412 [MꢀH]ꢀ (negative-ion mode) for the unaltered chelate.
General procedure for the reactions with potassium-graphite (C8K) lami-
nate:[31] In a two-neck round-bottom flask equipped with reflux condens-
er and magnetic stirring bar, graphite was heated (while stirred) under
argon for 15 min at 150–1608C. Potassium was added under argon, and
the mixture was kept at 1608C with careful stirring until the laminate
had formed (10–15 min). CAUTION: The material was highly pyrophor-
ic, necessitating cautious handling in thoroughly dried solvents. The dis-
tinctive bronze color of the mixture indicated that C8K was formed and
this was then suspended in anhydrous THF. A suspension of the chelate
Na(Fe-EDDHA) or the free ligand (o,o-EDDHA) in THF at room tem-
perature was added to this suspension. The ratio complex (or ligand)/
graphite/potassium was 1:48:6. This suspension was allowed to stir over-
night, quenched with water, and filtered through celite. The phases were
separated with CH2Cl2 and dried under vacuum. The aqueous layer was
analyzed. In the experiments with the free ligand (EDDHA), this prod-
uct was recovered unaltered as shown by NMR spectroscopy. In the ex-
[15] R. L. Chaney, J. Plant Nutr. 1984, 7, 47–67.
[16] P. R. Moog, W. Brꢁgermann, Plant Soil 1994, 165, 241–260.
[17] a) M. A. Sierra, M. J. MancheÇo, R. Vicente, M. Gómez-Gallego, J.
Org. Chem. 2001, 66, 8920–8925; b) E. Alvaro, M. C. de la Torre,
M. A. Sierra, Org. Lett. 2003, 5, 2381–2384; c) M. Gómez-Gallego,
R. Alcµzar, M. A. Sierra, F. Yunta, S. García-Marco, J. J. Lucena,
Dalton Trans. 2004, 3741–3747.
[18] a) M. Gómez-Gallego, M. J. MancheÇo, P. Ramírez-López, M. A.
Sierra, Tetrahedron 2000, 56, 4893–4905; b) M. J. MancheÇo, P.
Ramírez-López, M. Gómez-Gallego, M. A. Sierra, Organometallics
2002, 21, 989–992; c) M. A. Sierra, I. Fernµndez, M. J. MancheÇo,
M. Gómez-Gallego, M. R. Torres, F. P. Cossio, A. Arrieta, B. Lecea,
A. Poveda, J. Jimꢂnez-Barbero, J. Am. Chem. Soc. 2003, 125, 9572–
9573; d) P. Ramírez-López, M. Gómez-Gallego, M. J. MancheÇo,
M. A. Sierra, J. Org. Chem. 2003, 68, 3538–3545; e) P. Ramírez-
López, M. A. Sierra, M. Gómez-Gallego, M. J. MancheÇo, H. Gor-
nitzka, Organometallics 2003, 22, 5092–5099; f) M. A. Sierra, M. J.
MancheÇo, J. C. del Amo, I. Fernµndez, M. Gómez-Gallego, Chem.
Eur. J. 2003, 9, 4943–4953; g) R. Martínez-Alvarez, M. Gómez-Gal-
lego, I. Fernµndez, M. J. MancheÇo, M. A. Sierra, Organometallics
6004
www.chemeurj.org
ꢀ 2005 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Chem. Eur. J. 2005, 11, 5997 – 6005

Photochemistry of FeIII–EDDHA
FULL PAPER
2004, 23, 4647–4654; h) M. Gómez-Gallego, M. J. MancheÇo, M. A.
Sierra, Acc. Chem. Res. 2005, 38, 44–53; i) I. Fernµndez, M. A.
Sierra, M. J. MancheÇo, M. Gómez-Gallego, F. P. Cossio, Chem. Eur.
J. 2005, 11: DOI: 10.1002/chem.200400944.
Waltz, E. Zinato, D. W. Watts, P. D. Fleischauer, R. D. Lindholm,
Chem. Rev. 1968, 68, 541–585.
[29] a) M. Y. Darensbourg, S. Slater, J. Am. Chem. Soc. 1981, 103, 5914–
5915; b) J. M. Maher, R. P. Beatty, N. Cooper, Organometallics 1985,
4, 1354–1361.
[30] P. Girard, J. L. Namy, H. B. Kagan, J. Am. Chem. Soc. 1980, 102,
2693–2698.
[31] M. A. Sierra, P. Ramírez-López, M. Gómez-Gallego, T. Lejon, M. J.
MancheÇo, Angew. Chem. 2002, 114, 3592–3595; Angew. Chem. Int.
Ed. 2002, 41, 3442–3445.
[32] W. L. Lyndsay, Chemical Equilibria in Soils, Wiley, New York, 1979,
pp. 242–243.
[33] J. L. Pierre, M. Fontecave, R. R. Crichton, BioMetals 2002, 15, 341–
346.
[34] S. Dhungana, C. Ratledge, A. V. Crumbliss, Inorg. Chem. 2004, 43,
6274–6283.
[35] D. M. Miller, G. R. Buettner, S. D. Aust, Free Radical Biol. Med.
1990, 8, 95–108.
[19] B. C. Faust, R. G. Zepp, Environ. Sci. Technol. 1993, 27, 2517–2522.
[20] Plants have developed different methods to acquire iron from soils.
Strategy-I plants reduce exogenous FeIII from extracellular ferric
chelates, whereas strategy-II plants release chelators called phytosi-
derophores to take FeIII from the soil. The FeIII–phytosiderophore
complexes are taken back into the cell where the reduction occurs.
See also: D. Staiger, Angew. Chem. 2002, 114, 2363–2368; Angew .
Chem. Int. Ed. 2002, 41, 2259–2264.
[21] To be effective as soil fertilizer FeIII–EDDHA requires yearly appli-
cations either as solid or in aqueous solution. The excess of the
product remains absorbed in the soil where the effect of sunlight is
negligible. However, its final fate is to reach natural waters were it
would be exposed to sunlight and where the photodegradation
could be significant.
[22] A. Gilbert, J. Baggott, Essentials of Molecular Photochemistry,
Blackwell Science, London, 1991, pp 411–412.
[23] a) U. C. Yoon, Z. Su, P. S. Mariano in CRC Handbook of Organic
Photochemistry and Photobiology, 2nd ed. (Eds.: W. Horspool, F.
Lenci), CRC Press, Boca Raton, 2004, Chapter 101. b) J. D. Coyle,
Chem. Rev. 1978, 78, 97–123.
[24] F. Yunta, S. García-Marco, J. J. Lucena, M. Gómez-Gallego, R.
Alcµzar, M. A. Sierra, Inorg. Chem. 2003, 42, 5412–5421.
[25] a) G. D. Cooper, B. A. DeGraff, J. Phys. Chem. 1972, 76, 2618–
2625; b) M. Izakovic, J. Sima, V. Brezova, J. Photochem. Photobiol.
A 2004, 167, 81–86.
[36] C. J. Bannochie, A. E. Martell, J. Am. Chem. Soc. 1989, 111, 4735–
4742.
[37] Apoproteins and porphyrin have been proposed as ligands Y to ex-
plain the biological reduction of well-known FeIII–ferrisiderophore
complexes such as those of ferrioxamine B, iron–transferrin, and fer-
rienterobactin, which have reduction potentials in the range of
ꢀ0.450 V to ꢀ0.750 V. See: J. L. Pierre, M. Fontecave, Biometals
1999, 12, 195–199. See also reference [33].
[38] a) F. Yunta, S. García-Marco, J. J. Lucena, J. Agric. Food Chem.
2003, 51, 5391–5399; b) M. Gómez-Gallego, M. A. Sierra, R.
Alcµzar, P. Ramírez, C. PiÇar, M. J. MancheÇo, S. García-Marco, F.
Yunta, J. J. Lucena, J. Agric. Food Chem. 2002, 50, 6395–6399.
[39] M. A. Sierra, M. Gómez-Gallego, R. Alcµzar, J. J. Lucena, A. ꢃlvar-
ez, F. Yunta-Mezquita, Patent WO02/00604, January, 2002.
[40] N. A. Bailey, D. Cummins, E. D. McKenzie, J. M. Worthington,
Inorg. Chim. Acta 1981, 50, 111–120.
[26] K. Barbeau, G. Zhang, D. H. Live, A. Butler, J. Am. Chem. Soc.
2002, 124, 378–379.
[27] a) S. Fukuzumi in Advances in Electron-Transfer Chemistry, (Ed.:
P. S. Mariano), JAI Press, Greenwich, CT, 1992, pp. 67–175; b) S.
Fukuzumi, T. Tanaka in Photoinduced Electron Transfer (Eds.:
M. A. Fox, M. Chanon), Elsevier, Amsterdam, 1998, Part C, Chap-
ter 10.
[41] M. A. Cremonini, A. ꢃlvarez-Fernµndez, J. J. Lucena, A. Rombola,
B. Marangoni, G. Placucci, J. Agric. Food Chem. 2001, 49, 3527–
3532.
[28] a) V. Balzani, V. Carassiti, Photochemistry of Coordination Com-
pounds, Academic Press, London, 1970; b) A. W. Adamson, W. L.
Received: March 14, 2005
Published online: July 29, 2005
Chem. Eur. J. 2005, 11, 5997 – 6005
ꢀ 2005 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
www.chemeurj.org
6005