Iron(III) complexation by new aminocarboxylate chelators - Thermodynamic and kinetic studies

Article abstract of DOI:10.1002/1099-0682(200102)2001:2<471::AID-EJIC471>3.0.CO;2-A

The complexation of Fe^III by new tetradentate and pentadentate aminocarboxylate chelators, designed to protect cells against iron-catalysed oxidative damage, was investigated. Ferric iron complexes of N,N′-bis(3,4,5-trimethoxybenzyl)-ethylenediamine-N,N′ -diacetic acid (L1), N,N′-dibenzylethylenediamine-N,N′-diacetic acid (L2) and [N-(2-hydroxybenzyl)-N′-benzylethylenediamine-N,N′-di acetic acid] (L3) have been characterized in aqueous solution by potentiometric, UV/Vis spectrophotometric and cyclic voltammetric measurements. The parent ligand, ethylenediamine-N, N′-diacetic acid (L4), has also been studied. As expected, the presence of a hard phenolate donor group in L3 significantly enhances the affinity for iron while decreasing the potential of the Fe^III/Fe^II redox couple compared to L1 or L2. Kinetics studies have provided the kinetic rate constants related to the formation and the dissociation of the ferric complex with L3. The results reveal a fast Fe^III uptake, which is a favorable feature for a biological use of this type of ligand. Overall, these results demonstrate the pertinence of the use of such ligands to protect biological tissues against oxidative stress.

Full text of DOI:10.1002/1099-0682(200102)2001:2<471::AID-EJIC471>3.0.CO;2-A

FULL PAPER
Iron(III) Complexation by New Aminocarboxylate Chelators ؊
Thermodynamic and Kinetic Studies
Guy Serratrice,*[a] Jean-Baptiste Galey,[c] Eric Saint Aman,[b] and Jacqueline Dumats[c]
Keywords: Iron / Chelates / Thermodynamics / Kinetics / N ligands
The complexation of FeIII by new tetradentate and penta-
dentate aminocarboxylate chelators, designed to protect cells
against iron-catalysed oxidative damage, was investigated.
Ferric iron complexes of N,NЈ-bis(3,4,5-trimethoxybenzyl)-
ethylenediamine-N,NЈ-diacetic acid (L1), N,NЈ-dibenzylethy-
acid (L4), has also been studied. As expected, the presence
of a hard phenolate donor group in L3 significantly enhances
the affinity for iron while decreasing the potential of the FeIII/
FeII redox couple compared to L1 or L2. Kinetics studies have
provided the kinetic rate constants related to the formation
lenediamine-N,NЈ-diacetic acid (L2) and [N-(2-hydroxyben- and the dissociation of the ferric complex with L3. The results
zyl)-NЈ-benzylethylenediamine-N,NЈ-diacetic acid] (L3) have
reveal a fast FeIII uptake, which is a favorable feature for a
been characterized in aqueous solution by potentiometric, biological use of this type of ligand. Overall, these results
UV/Vis spectrophotometric and cyclic voltammetric meas- demonstrate the pertinence of the use of such ligands to pro-
urements. The parent ligand, ethylenediamine-N,NЈ-diacetic
tect biological tissues against oxidative stress.
Introduction
their normal storage sites and become available to catalyze
Fenton chemistry, leading to oxidative damage to sur-
rounding biological molecules.[5] However, there are few re-
ports on the development of iron chelators specifically de-
signed for the treatment of such oxidative stress-associated
situations. The main reason is probably that the design of
such systems is potentially problematic for medicinal chem-
ists in terms of safety margins owing to the possible interac-
tion of chelators with normal iron metabolism.[6]
In recent years there has been an intensive search for new
iron ligands with medicinal applications. Several diseases
are indeed associated with disturbance of iron homeost-
asis.[1] One major issue is the treatment of iron-overload-
associated situations for people suffering from haematolog-
ical defects such as β-thalassemia major. Indeed, because
man is unable to actively excrete iron, systemic iron over-
load occurs rapidly when frequent blood transfusions are
given.[2] In such situations, treatment with powerful iron
chelators is the only effective way to remove excess iron.
Currently, the hexadentate siderophore desferrioxamine
(DFO) is by far the clinically most used iron chelator. How-
ever, DFO lacks oral activity and has a very short biological
half life, which results in poor patient compliance.[3] There-
fore, considerable effort has been focused on discovering
new therapeutically useful iron chelators.
On the other hand, there are also other areas in which
iron chelators could be useful. For instance, drugs able to
specifically chelate iron from iron-containing enzymes such
as ribonucleotide reductase or lipoxygenase could find im-
portant applications.[4] Moreover, in addition to systemic
iron overload, local iron homeostasis disturbances can oc-
cur in several conditions associated with so-called ‘‘oxidat-
ive stress’’, i.e. when traces of iron are locally released from
We recently developed a new strategy based on the idea
that an ‘‘ideal’’ iron chelator for use in oxidative stress-re-
lated conditions should effectively chelate iron only in pro-
oxidant conditions. We therefore explored the possibility of
generating pro-drugs that can be oxidatively activated into
species with strong iron-binding capacity.[7] Of particular
interest in such a strategy is the degree of control exercised
by the oxidative stress conditions, whereby a dormant rela-
tively inactive compound is transformed to an active species
by the very conditions prevailing at sites where it is most
needed.
N,NЈ-Bis(3,4,5-trimethoxybenzyl)ethylenediamine-N,NЈ-
diacetic acid (L1) (Figure 1) is the most extensively studied
ligand in this series. In aqueous solutions containing acetate
and ferric salts, L1 forms a µ-acetato µ-oxo diferric iron
complex.[8] In the presence of ferrous salts in anaerobic con-
ditions, a mononuclear ferrous complex is formed. Both
complexes react with hydrogen peroxide or dioxygen in the
presence of ascorbate to form a mononuclear (phenolato)-
FeIII complex resulting from the quantitative hydroxylation
of L1 into L5 (Figure 2). Owing to the well-known affinity
of phenolate ligands for ferric iron, L1 appears to fit the
aforementioned prerequisites regarding the generation of a
strong iron chelator according to a local oxidative activa-
tion process.[9]
[a]
´
Laboratoire de Chimie Biomimetique, LEDSS, UMR CNRS
´
5616, Universite J. Fourier
B. P. 53, 38041 Grenoble, Cedex 9, France
E-mail: Guy.Serratrice@ujf-grenoble.fr
[b]
[c]
Laboratoire dЈElectrochimie Organique et de Photochimie
´
Redox, UMR CNRS 5630, Universite J. Fourier,
B. P. 53, 38041 Grenoble, Cedex 9, France
´
´
´
Departement Molecules & Materiaux, l’OREAL Recherche
´
Avancee,
1 Avenue Eugene Schueller, 93600 Aulnay sous Bois, France
`
Eur. J. Inorg. Chem. 2001, 471Ϫ479
WILEY-VCH Verlag GmbH, 69451 Weinheim, 2001
1434Ϫ1948/01/0202Ϫ0471 $ 17.50ϩ.50/0
471

G. Serratrice, J.-B. Galey, E. Saint Aman, J. Dumats
FULL PAPER
ure 1), i.e. stability constants and redox potentials to dem-
onstrate the pertinence of the use of such ligands to protect
biological tissues against iron-catalyzed oxidative damage.
The ligand ethylenediamine-N,NЈ-diacetic acid (L4) was
also investigated for comparison and because the stability
constants of the ferric complexes are unknown. A detailed
kinetic study of the formation and of the dissociation of
the ferric complex with L3 is also described in order to (i)
determine the reactivity patterns of this complex containing
phenolate and aminoacetate donor groups and (ii) since few
data are available in the literature for this type of ligand.
Results and Discussion
Ligand Deprotonation Constants
The deprotonation constants Ka of the ligands investig-
ated in this research were studied by potentiometric titra-
tion. Analysis of the potentiometric titration curve (Fig-
ure 3a and Figure 3c as examples) by the SUPERQUAD
program[10] yielded the pKan values defined by Equation (1)
and Equation (2) (charges omitted for clarity) and reported
in Table 1.
Kan ϭ [LHn Ϫ 1]·[Hϩ]/[LHn]
(2)
Figure 1. Structure of the ligands Ln
Three pKa values were determined and these correspond
to one carboxylic acid group and the two ammonium
groups. By comparison to EDTA the lowest pKa value for
each ligand is attributed to a carboxylic acid and the two
highest ones to the ammonium nitrogen atoms. The pKa
values of L4 are in good agreement with those determined
by Harris and Martell [9.67, 6.57 and 2.36 (I ϭ 0.1 ,
KNO3)].[11] The pKa value of the other carboxylic acid is
assumed to be significantly lower than 2.
A fourth pKa value corresponding to the hydroxy group
was determined for the ligand L3 by spectrophotometric
titration. This pKa cannot be accurately determined by po-
tentiometric titrations since deprotonation occurs at pH Ͼ
11. A spectrophotometric titration was carried out over the
pH range 10Ϫ12. The corresponding UV spectra exhibit an
isosbestic point at λ ϭ 270 nm, indicating the presence of
only two absorbing species. The data were processed by the
LETAGROP SPEFO[12,13] program (absorbance values at 8
wavelengths). The best fit [Σ(Aexp Ϫ Acalc)2 ϭ 10Ϫ3] yielded
a pKa value of 12.0 Ϯ 0.1. The extinction coefficients were
determined to be εmax ϭ 3500 Ϫ1cmϪ1 (λmax ϭ 292 nm)
and εmax ϭ 10000 Ϫ1cmϪ1 (λmax ϭ 238 nm) for L33Ϫ and
εmax ϭ 50 (λmax ϭ 292 nm) and εmax ϭ 1000 Ϫ1cmϪ1
(λmax ϭ 238 nm) for L3H2Ϫ. The pKa values for L3 are
Figure 2. Reaction scheme for the oxidation of Fe2(L1)2
In the present paper, we extend our previous findings by close to those obtained for the ligands HBIDA [N-(2-hydro-
accurately investigating the physicochemical parameters of xybenzyl)iminodiacetic acid][14] and HBED [N,NЈ-bis(2-hy-
FeIII complexation by L1 and analogs L2 (N,NЈ-dibenzyle- droxybenzyl)ethylenediamine-N,NЈ-diacetic acid][15] bearing
thylenediamine-N,NЈ-diacetic acid) and L3 [N-(2-hydroxyb- similar donor groups, as indicated in Table 1. The high basi-
enzyl)-NЈ-benzylethylenediamine-N,NЈ-diacetic acid] (Fig- city of the hydroxy group reflects the presence of a hydro-
472
Eur. J. Inorg. Chem. 2001, 471Ϫ479

Iron(III) Complexation by New Aminocarboxylate Chelators
FULL PAPER
a ϭ 5, where a is equal to the number of moles of base
added per mole of ligand. This clearly indicates the forma-
tion of hydroxo complexes Fe(OH)L and Fe(OH)2L at rela-
tively low pH values. Analysis of the titration curves by the
SUPERQUAD program[10] yielded the β110 and Kn (n ϭ
OH
1 or 2) constants expressed as in Equation (3) and Equation
(4).The values are given in Table 2.
β110 ϭ [FeL]/[Fe3ϩ]·[L]
Kn ϭ [Fe(OH)nL]·[Hϩ]/[Fe(OH)n Ϫ 1L]
(3)
(4)
OH
Table 2. FeIIIϪligand stability constants and pFe values
Constants[a]
L1
L2
L3
L4
log β110
13.95(4)
3.34(4)
7.34(5)
17.2
15.20(9)
3.52(9)
7.27(5)
18.1
27.00(8)[b]
5.76(8)[b]
15.50(1)
3.64(5)
7.18(2)
18.2
[c]
pK1
pK2OH
[d]
OH
pFe[e]
23.3
[a]
All values were determined at 25 °C and I ϭ 0.1  (NaClO4);
values in parentheses are standard deviations in the last significant
digit. Ϫ [b] Determined from spectrophotometric titration; K1
ϭ
ϭ [Fe(OH)L]O[HHϩ]/
[c]
6.02 from potentiometric titration. Ϫ
[FeL] for the equilibrium FeL ϩ H2O
K1
OH
Fe(OH)L ϩ Hϩ. Ϫ
Ǟ
[d]
K2OH ϭ [Fe(OH)2L][Hϩ]/[Fe(OH)L] forǟthe equilibrium FeLOH ϩ
Fe(OH)2L ϩ Hϩ. Ϫ
Calculated for [L]tot ϭ 10Ϫ5 ,
Ǟ
[e]
H2O
ǟ
[Fe]tot ϭ 10Ϫ6  at pH ϭ 7.4.
For the tetraprotonated ligand L3, the breaks at a ϭ 4
and a ϭ 5 indicate the formation of the FeL and Fe(OH)L
complexes (Figure 3b). However, it was not possible to de-
termine the β110 value since the formation of the complex
is complete at low pH values. Only the value of K1
was
Figure 3. Potentiometric titration curves for (a) 1 m ligand L3;
(b) L3 ϩ Fe3ϩ 1:1, 1 m; (c) 1 m ligand L1; (d) L1 ϩ Fe3ϩ 1:1,
1 m; a ϭ mol of base added per mol of ligand; all solutions were
at 25 °C and I ϭ 0.1  (NaClO4); the data were refined by the
SUPERQUAD program (σfit ϭ 3.5Ϫ4.5)
OH
determined from titration data and this is given in Table 2.
The spectrophotometric titration was carried out over the
range pH ϭ 1Ϫ8 in order to determine the constants β110
and K1OH. The stoichiometric addition of FeIII to a solution
of L3 over the pH range 1Ϫ2 led to the formation of a
purple complex, the UV/Vis spectrum of which shows a
maximum at 525 nm that is attributed to a phenolato-to-
FeIII charge transfer transition (Figure 4a). The absorbance
data were processed with the LETAGROP-SPEFO pro-
gram.[12,13] The best fit [Σ(Aexp Ϫ Acalc)2 ϭ 2 ϫ 10Ϫ3] was
obtained by considering the formation of the [FeL] species
Table 1. Deprotonation constants of the ligands
[a]
pKan
L1
L2
L3
L4
HBIDA[b] HBED[c]
pKa1
pKa2
pKa3
pKa4
pKa5
9.44(1) 9.68(1) 12.0(1) 9.75(1) 11.71
4.61(2) 4.93(1) 9.11(1) 6.68(2) 8.07
2.87(2) 2.29(3) 4.67(1) 2.64(3) 2.34
2.17(2)
12.35
12.08
8.46
4.76
2.18
and provides the values log β110 ϭ 27.0 Ϯ 0.08 and εmax
ϭ
[a]
All values were determined at 25 °C and I ϭ 0.1  (NaClO4);
1950 Ϫ1cmϪ1 (λmax ϭ 520 nm). No spectral change was
observed over the pH range 2.5Ϫ4.5. As the pH value was
increased from 5 to 7, the spectra showed isosbestic points
at 394 and 491 nm, indicating that there are only two spe-
cies with a different absorption in this pH range (Fig-
ure 4b). These absorbance data were also refined with the
values in parentheses are standard deviations in the last significant
[b]
[c]
digit. Ϫ
N-(2-Hydroxybenzyl)iminodiacetic acid, ref.[14]
Ϫ
N,NЈ-Bis(2-hydroxybenzyl)ethylenediamine-N,NЈ-diacetic
acid,
ref.[15]
gen bond between the amino nitrogen atom and the hy-
droxy group.
LETAGROP-SPEFO program.[12,13] The best fit [Σ(Aexp
Ϫ
2
Acalc
)
ϭ 3 ϫ 10Ϫ3] showed that deprotonation occurs
Stability Constants of the Ferric Complexes
through a one-proton step yielding a pK1
value of 5.76
OH
The stability constants β110 were determined by potentio- Ϯ 0.08 and an absorption maximum at λ ϭ 470 nm with
metric titration for the ferric complexes of ligands L1, L2 ε ϭ 1800 Ϫ1cmϪ1 for the Fe(OH)L species. The value of
and L4. The titration curves for a 1:1 ratio of FeIII to ligand pK1OH is in good agreement with the value 6.02 determined
(Figure 3d as an example) exhibit two breaks at a ϭ 4 and by potentiometric titration. The shift of the absorption
Eur. J. Inorg. Chem. 2001, 471Ϫ479
473

G. Serratrice, J.-B. Galey, E. Saint Aman, J. Dumats
FULL PAPER
Figure 5. Distribution diagrams of the system FeIIIϪL3 (a) and
FeIIIϪL1 (b) as a function of pH; [Fe3ϩ] ϭ 1 m, [ligand] ϭ 10 m
Figure 4. UV/Vis absorption spectra of Fe3ϩϪL3 as a function of
pH; (a) 1: pH ϭ 0.9; 2: pH ϭ 1.8; (b) 1: pH ϭ 4.5; 2: pH ϭ 6.5;
[Fe3ϩ] ϭ [L3] ϭ 0.5 m; I ϭ 0.1  (NaClO4)
agreement with those of tetradentate N,NЈ-diethylenediam-
ine diacetic acid derivatives. Consequently, low values for
pK1
are observed (over the range 3.2Ϫ3.6). This reflects
OH
maximum to lower wavelength reflects coordination with the high Lewis acid character of the ferric ion in FeL re-
the more basic hydroxide ion. A similar spectral change has sulting from the low basicity of these ligands. A more reli-
been observed for the ferric complex with HBIDA[14] able parameter for comparison of ligand effectiveness is the
[λmax ϭ 518 nm, εmax ϭ 1100 Ϫ1cmϪ1 for FeL, λmax
ϭ
pFeIII value (ϭ Ϫlog [Fe3ϩ]). The larger the pFe value, the
460 nm, εmax ϭ 1010 Ϫ1cmϪ1 for Fe(OH)L]. It is worth more effective the ligand. The pFeIII values were calculated
noting that the extinction coefficient is about half the value for [L]tot ϭ 10Ϫ5 , [Fe]tot ϭ 10Ϫ6  at pH ϭ 7.4 and are
for Fe-HBED[15] (λmax ϭ 485 nm, εmax ϭ 3935 Ϫ1cmϪ1), reported in Table 2. The pFe values for L2 and L4 are very
a system that has two phenolate donors and the same similar, indicating that substitution with benzyl groups does
framework as L3. The distribution curves are shown in Fig- not influence the complexation. The lower ferric affinity of
ure 5 for ligands L1 and L3.
L1 in comparison to L2 and L4 is probably due to steric
The ferric complexes of the tetradentate ligands L1, L2 effects of the methoxy substituents. A higher metal ion af-
and L4 exhibit relatively low stability constants β110, in finity is observed for L3 and this fact is related to the
474
Eur. J. Inorg. Chem. 2001, 471Ϫ479

Iron(III) Complexation by New Aminocarboxylate Chelators
FULL PAPER
phenolate donor, as one would expect. An increase in the results in the progressive appearance of a new redox peak
pFe value of 5.2 from L2 (18.1) to L3 (23.3) is observed. system at E1/2 ϭ Ϫ0.48 V (∆Ep ϭ 0.31 V), which grows at
A comparison with the hexadentate ligand HBED (pFe ϭ the expense of the original one at E1/2 ϭ Ϫ0.21 V (Figure 6,
26.97),[15] which contains two phenolate donors with an curves 2Ϫ4). Above pH ϭ 8, the original wave disappears
EDTA-type framework, shows that incorporation of one and the new peak system reaches full development (Fig-
phenolate donor increases the pFe value by about 4Ϫ5 un- ure 6, curves 5 and 6). This result is in agreement with the
its. Furthermore, the pFe value for L3 is higher than that pK1
value of 5.76 found for the FeL3/Fe(OH)L3 system
OH
of the tetradentate ligand HBIDA (19.95), which has one and the distribution curve in Figure 5 (see above), i.e. the
phenolate donor, one nitrogen donor and two carboxylate wave at E1/2 ϭ Ϫ0.48 V can be attributed to the
donors. The pK1 values are very close, being 5.76 for L3 FeIII(OH)L3/FeIIL3 redox couple. Therefore, at pH ϭ 7.4,
OH
and 5.75 for HBIDA. This indicates that the Lewis acidity the L3 ferric chelate appears to be mainly in a form that is
of FeIII is similar in both complexes.
reducible by physiological reductants. However, the small
difference between the E1/2 values of NADPH and the L3
ferric chelate could explain the fact that L3Ϫiron is not a
catalyst of reductant-driven Fenton reaction under physio-
logical conditions.[7] In addition, a small proportion of the
complex remains in the FeL3 form, which could be reduced
by physiological reductants. This process would allow po-
tential redox cycling of iron, especially during acidosis,
which is generally associated with cell oxidative stress.
Therefore, it is likely that additional parameters, such as the
kinetics of complex formation, compensate for the thermo-
dynamic factor that is favorable for the catalysis of Fenton
chemistry.
Electrochemistry
The electrochemical behavior of the FeIII ϩ L(1؊3) com-
plexes was studied by cyclic voltammetry in an aqueous so-
lution over the pH range 3Ϫ10.5 in order to investigate the
ability of these complexes to undergo redox cycling under
physiological conditions and to catalyze Fenton chemistry.
In order to be able to catalyze the formation of hydroxyl
radicals in the presence of hydrogen peroxide, an iron com-
plex has to meet two conditions simultaneously: (i) the fer-
ric chelate has to be reducible by physiological reductants,
i.e. its standard redox potential must be higher than Ϫ0.534
V vs. Ag/AgCl (NADPH/NADPϩ redox couple) and (ii)
single electron transfer from the ferrous chelate to hydrogen
peroxide must be possible, i.e. its redox potential must lie
below 0.25 V vs. Ag/AgCl (H2O2/HO·, OHϪ). The electro-
chemical response of the FeIII ϩ L3 electrolytic solution is
characterized in the 3Ϫ5 pH range by a quasi-reversible
redox wave at E1/2 ϭ Ϫ0.21 V vs. Ag/AgCl (∆Ep ϭ 0.12 V,
v ϭ 0.1 V sϪ1) corresponding to the FeIII/FeII redox couple
(Figure 6, curve 1). Increasing the pH value from 5 to 8
In contrast, the electrochemical behavior of L(1؊2) ϩ
FeIII complexes appears more complicated, possibly due to
the presence of several species in solution at basic pH
values, including the µ-oxo binuclear species already char-
acterized by ESI-MS.[8] At pH values above 8, an ill-be-
haved peak system is seen in the CV curves (Figure 7,
curves 1 and 2) (Epc ϭ Ϫ0.80 and Ϫ0.72 V, Epa ϭ Ϫ0.15
and Ϫ0.04 V for the complexes L2 ϩ FeIII and L1 ϩ FeIII
,
respectively). For pH values lower than 4.5 (Figure 7,
Figure 7. CV curves of FeIII ϩ L2 (1, 3) and FeIII ϩ L1 (2, 4),
0.5 m in H2O ϩ NaClO4 0.1  at a GC electrode (5 mm dia-
Figure 6. CV curves of FeIII ϩ L3 (2.9 m) in H2O ϩ NaClO4 0.1
 at a GC electrode (5 mm diameter); scan rate 0.1 VsϪ1; E vs. Ag/ meter); scan rate 0.1 V sϪ1; E vs. Ag/AgCl, aqueous NaCl (3 )
AgCl, aqueous NaCl (3 ) (ϩ 0.21 V vs. NHE); pH ϭ 4.2 (1), 5.4 (ϩ0.222 V vs. NHE); pH ϭ 8.3 (1, 2), 3.0 (3, 4) (adjusted with
(2), 6.3 (3), 7.0 (4), 8.3 (5), 9.2 (6) (adjusted with HClO4 or NaOH) HClO4 or NaOH); scale (S) ϭ 10 µA (1, 2), 4 µA (3, 4)
Eur. J. Inorg. Chem. 2001, 471Ϫ479
475

G. Serratrice, J.-B. Galey, E. Saint Aman, J. Dumats
FULL PAPER
curves 3 and 4), a reversible redox wave is seen at E1/2
ϭ
0.21 and 0.22 V (∆Ep ϭ 0.06 and 0.08 V, respectively) in
the CV curves recorded in the presence of L2 ϩ FeIII and
L1 ϩ FeIII, respectively. At lower potentials, a fully irrevers-
ible peak at 0 V (L2 ϩ FeIII) or Ϫ0.04 V (L1 ϩ FeIII) is
observed, confirming the presence in solution of two differ-
ent FeIII species in a frozen equilibrium, i.e. the interconver-
sion of these species is slow on the time scale of voltamme-
try. Between pH ϭ 5 and 8, the CV curves consist of the
overlapped electrochemical signals observed in basic and
acidic media. These results are in agreement with pK1
OH
and pK2 values of ca. 3 and 7 found for both L2 ϩ FeIII
OH
and L1 ϩ FeIII complexes, i.e. in basic media only the
FeIII(OH)2L complex is present in solution whereas in
acidic media an equilibrium between FeIIIL and FeIII(OH)L
occurs. Taking into account the negative shift in potential
for the redox couple FeIII(OH)L/FeIIL in comparison to
FeIIIL3/FeIIL3, it can be assumed that the reversible wave
at around 0.2 V is attributable to the FeIIIL/FeIIL redox
couple (where L signifies L1 or L2) when the irreversible
electron transfer corresponds to the cathodic process char-
acteristic of FeIII(OH)L3. The irreversibility of the electron
transfer is likely to be due to a change in the coordination
sphere around the Fe center upon reduction. Comparing
L1 and L2, it must be noted that the substitution of L1
with electron-donating methoxy groups does not influence
significantly the potential characteristic of the iron-local-
ized electron transfer. In addition, under the same experi-
mental conditions, E1/2 for the FeIIIEDTA/FeIIEDTA redox
couple is found at Ϫ0.10 V, showing that L3 stabilizes the
FeIII form of the complex by ca. 0.12 V while L1 and L2
stabilize the FeII form of the complex by ca. 0.31 V. This is
in agreement with the presence of the hard coordinating
hydroxo site on L3.
Figure 8. Variation of the experimental rate constant kobs (sϪ1) for
the formation of the FeIIIϪL3 complex as a function of [Fe3ϩ] ()
at various [Hϩ] (): (᭜) 0.05, (ᮀ) 0.08, (᭹) 0.10, (∆) 0.125, (᭡)
0.15; [L3] ϭ 0.4 m; solvent: water, I ϭ 2.0  (HClO4 ϩ NaClO4),
T ϭ 25 °C
The acid hydrolysis kinetics of the FeϪL3 complex was
studied in aqueous solution (ionic strength I ϭ 2 , 25 °C)
by the pH jump method. The pH jump reactions were per-
formed under pseudo-first-order conditions with respect to
[Hϩ] (varying over the range 0.02Ϫ0.9 ) for solutions con-
taining FeIII and L3 in a 1:1 molar ratio at an initial pH
value of 4. The UV/Vis spectra recorded as a function of
the time resemble the spectra recorded for the formation
reaction study and for the equilibrium study at pH Ͻ 2
(Figure 4a) and indicate a total dechelation at the iron
center. The exponential absorbance decay vs. time recorded
at 520 nm indicates that the dissociation occurs through a
single determining step. The plot of the kobs rate constants
vs. [Hϩ] is linear over the [Hϩ] range 0.2Ϫ1.0  and slightly
curved at low [Hϩ] (Figure 9).
Formation and Acid Hydrolysis Kinetics of the FeL3
Complex
The capacity of a chelator to protect against the forma-
tion of damaging oxidizing species is influenced not only by
several thermodynamic factors, but also by kinetic factors.
Therefore, the kinetics of formation of the FeL3 complex
was investigated by a stopped-flow spectrophotometric
method under pseudo-first-order conditions: [FeIII] and
[Hϩ] ϾϾ [L] and at ionic strength I ϭ 2 , 25 °C. The
absorbance change vs. time of the phenolato-to-FeIII charge
transfer band at λ ϭ 525 nm showed a single exponential
curve indicating that the complex formation reaction pro-
ceeds through a single rate-limiting step. The spectra re-
corded with time using a diode array device were similar to
those shown in Figure 4a. It was found that the pseudo
first-order rate constants kobs at a given acidity level ([Hϩ]
varied from 0.05 to 0.15 ) have a linear variation as a
function of [Fe3ϩ] with a significant intercept (Figure 8)
that indicates a contribution of the reverse reaction. The
slopes of these plots decrease as [Hϩ] increases. In addition,
the intercepts of the plots increase as [Hϩ] increases.
Figure 9. kobs (sϪ1) as a function of [Hϩ] () in the acid hydrolysis
kinetics of the FeIIIϪL3 complex, [FeL3] ϭ 0.5 m; solvent: water;
I ϭ 2.0  (HClO4 ϩ NaClO4), T ϭ 25 °C
476
Eur. J. Inorg. Chem. 2001, 471Ϫ479

Iron(III) Complexation by New Aminocarboxylate Chelators
By taking into account the equilibria existing in solution
FULL PAPER
(kϪ3 ϩ kϪ4) ϭ 1.83 Ϯ 0.09 Ϫ1sϪ1 kϪ5 ϭ 0.42 Ϯ 0.08
under the experimental conditions and after analyzing the sϪ1 KFeLH2 ϭ 0.026 Ϯ 0.005 2
kinetics results from both the formation and the hydrolysis
The rate constants of the Fe3ϩ pathways have been found
reactions, the reaction scheme shown in Equation (5) to to be generally almost three orders of magnitude lower than
Equation (14) can be proposed ([FeOH]2ϩ for those of the [FeOH]2ϩ pathways.[18] So, if we assume a value
[Fe(OH)(OH2)5]2ϩ, Fe3ϩ for [Fe(OH2)6]3ϩ and L for L3).
of about 10 Ϫ1sϪ1 for k4, it seems reasonable to neglect
the k4Ka4Ka5 term with respect to k3KFeKa5. Thus, we can
obtain k3 ϭ 770 Ϯ 180 Ϫ1sϪ1. From the values of k5 and
k
Ϫ5, k4/kϪ4 ϭ KFe k5/kϪ5 is calculated to be 47.5. An upper
limit of kϪ4 is thus estimated to be 0.2 if k4 Ͻ 10 and so
we can obtain kϪ3 ϭ 1.6 sϪ1
.
In summary, we were able to determine the following
rate constants:
k5 ϭ13300 Ϯ 1400 Ϫ1sϪ1; kϪ5 ϭ 0.42 Ϯ 0.08 sϪ1
k3 ϭ 770 Ϯ 180 Ϫ1sϪ1; kϪ3 ϭ 1.6 Ϯ 0.1 sϪ1
Ligand substitution reactions at an iron center were as-
sumed to be controlled by water exchange from the inner
coordination shell. The rate constants may be described ac-
cording to the EigenϪWilkins mechanism[19] with a fast
formation step of an outer-sphere complex (Kos) followed
by a rate-limiting step involving the water substitution of
the [FeOH]2ϩ species (kex) according to Equation (16); kf
stands for k3 or k5, S is a statistical coefficient that accounts
for the solvent shell composition and can be assigned an
approximate value of 1/6.
kf ϭ Kos·S·kex
(16)
Since kex does not change significantly with different li-
gands, variations in the rate constants are mainly due to the
constant Kos. The charge of the ligand is one of the factors
that influences the value of Kos. The constant can be estim-
ated with the Fuoss equation[20] to be 0.2 for reactions of
[FeOH]2ϩ with a neutral ligand and ca. 0.02 for a ligand
bearing a global positive charge. Accordingly, the value of
k5, which refers to the [FeOH]2ϩ Ϫ LH3 pathway, is larger
It was assumed that the rate-determining pathway in-
volves the formation of the [FeLH23ϩ] complex since the
reaction proceeds through the coordination of Fe3ϩ with
the phenolato oxygen atom and releases only one Hϩ in the
reaction according to Equation (12).
The pseudo-first-order kinetic according to Equa-
tion (15) can be obtained by neglecting terms from reac-
tions according to Equation (8) to Equation (10), in agree-
ϩ
than the value of k3, which refers to the [FeOH]2ϩ Ϫ LH4
ment with the experimental results, where [Fe3ϩ] ϭ [Fe3ϩ
]
tot
pathway (k5/k3 ഠ 17). Furthermore, a comparison with the
literature data indicates that k5 is much higher than the rate
constants determined for the [FeOH]2ϩ Ϫ neutral ligand
pathway for phenol, catechol and hydroxamate ligands
(about 1 ϫ 103 to 3 ϫ 103 Ϫ1sϪ1). This finding can be
explained by taking into account the zwitterionic character
of the ligand L3. The two carboxylate anions are assumed
to increase the Kos constant with respect to a ligand bearing
no charges. The values of k5 fall within the range for a li-
gand containing one or two aminoacetate donor groups,[21]
such as H4EDTA (30000 Ϫ1sϪ1) and H3NTA (15000
Ϫ [FeOH2ϩ] Ϫ 2 [Fe2(OH)24ϩ] and [Hϩ] takes into account
the protons involved in formation of [FeOH]2ϩ and
[Fe2(OH)2]4ϩ (2 Fe3ϩ Ǟ [Fe2(OH)2]4ϩ ϩ 2 Hϩ, KDFe).
ǟ
The values of KFe and KDFe in 2.0  NaClO4 are 1.5 ϫ
10Ϫ3  and 2.4 ϫ 10Ϫ3 M.[16] Since pKa5 was not deter-
mined, the fit of the data was performed using the value
1.17, which was obtained for the corresponding depro-
tonation of HBED.[15] The variation of the pseudo-first-or-
der kinetic rate constant of the hydrolysis reaction is con-
sistent with the last term of Equation 15.[17] The results of
the nonlinear least-squares fit to Equation (15) of the com-
bined formation and dissociation data are:

Ϫ1sϪ1).
Kinetic data for the dissociation of ferric complexes with
this type of ligand are not available in the literature. A com-
parison can be made with dissociation rate constants ob-
tained for catecholate and hydroxamate ligands, which are
the most widely reported in the literature. The dissociation
rate of FeL3 by the [FeOH]2ϩ pathway (1.6 and 0.42 sϪ1
)
is close to that of the (monocatecholato)FeϪTiron (1,2-di-
(k3KFeKa5 ϩ k4Ka4Ka5) ϭ 0.078 Ϯ 0.018 Ϫ1sϪ1 k5
13300 Ϯ 1400 Ϫ1sϪ1
ϭ
hydroxy-3,5-benzenedisulfonate) complex (1.2 sϪ1 [22]
)
but
significantly faster than those of the monohydroxamato
Eur. J. Inorg. Chem. 2001, 471Ϫ479
477

G. Serratrice, J.-B. Galey, E. Saint Aman, J. Dumats
Synthesis of L3
FULL PAPER
complexes (0.08 sϪ1 for acetohydroxamic acid and 0.0071
s
Ϫ1 for N-methylacetohydroxamic acid).[23,24] A comparison
(a) N-(2-Hydroxybenzyl)-NЈ-benzylethylenediamine: A mixture of
8.05 g (66 mmol) of 2-hydroxybenzaldehyde and 10 g (66.6 mmol)
of N-benzylethylenediamine was heated in 100 mL of MeOH at
50 °C for 30 min. The mixture was concentrated, the residue sus-
pended in 150 mL of EtOH and 34 mmol of NaBH4 was added.
The reaction mixture was stirred at room temperature for 1 h and
the solvent evaporated. 100 mL of water was added and concen-
trated HCl was used to acidify the solution to pH ϭ 1Ϫ2. The
crystalline dihydrochloride was collected by filtration, washed with
cold ethanol and vacuum-dried over P2O5 to yield 13.85 g (81%)
of N-(2-hydroxybenzyl)-NЈ-benzylethylenediamine. Ϫ MS (ESIϩ);
m/z: 256 [M ϩ H]ϩ, 91 [C7H7ϩ].
with the last step of the dissociation of hexadentate chelates
shows that L3 releases FeIII much faster than the tricatech-
olate TRENCAMS[25] (0.018 sϪ1) and than the sideroph-
ores pyoverdin[26] (Ͼ 6 10Ϫ5 Ϫ1) and ferrioxamine B[27,28]
s
(0.0021 sϪ1).
Conclusion
In summary, this work has shown that the bis(aminocar-
boxylate)benzyl chelators of the present series have a rela-
tively low affinity for FeIII. It must be emphasized that these
chelators (except L3) should be thermodynamically unable
to compete for iron in most metalloproteins and especially
for transferrin. For this reason side-effects can be expected
to be limited, a situation in contrast to strong iron
chelators, which can inhibit enzymes by removal of metal
from active sites, by formation of a ternary complex or by
depriving the apoenzyme of its normal source of iron. Un-
der prooxidant conditions, e.g. in the presence of hydrogen
peroxide, these chelators are hydroxylated to give species
with strong iron affinity, as shown by the complexation
stability constants and pFe value for L3, which are consist-
ent with the presence of a hard phenolate donor group.
Moreover, this dramatic increase of affinity for iron is asso-
ciated with a large decrease in redox potential of FeIII/FeII,
which tends to inhibit iron reduction by physiological re-
ductants thereby avoiding redox cycling of iron. However,
it should be kept in mind that such a consideration refers
to equilibrium conditions, which is generally not the case in
biological systems, and these results agree with previous
data showing that the L3 iron complex is not a catalyst of
reductant-driven Fenton reactions.[7] In addition, both the
formation and dissociation kinetics of the ferric complex
with L3 were measured and were found to give consistent
data. It should be emphasized that an interesting feature of
L3 is the fast uptake of FeIII and its fast release under acidic
conditions. The results presented here are in good agree-
ment with biological data showing a very efficient protec-
tion of cultured cells by L3 or L1 prodrugs against hydro-
gen peroxide toxicity.[29,9]
(b) N-(2-Hydroxybenzyl)-NЈ-benzylethylenediamine Diacetic Acid:
A mixture of 12.8 g (50 mmol) of N-(2-hydroxybenzyl)-NЈ-benzyl-
ethylenediamine, 13.9 g (100 mmol) of bromoacetic acid, 2 g
(50 mmol) of NaOH and 8.4 g (100 mmol) of NaHCO3 in 150 mL
of water was heated at 40 °C for 6 h while the pH value of the
solution was maintained in the range of 11.5Ϫ12.5 by the addition
of NaOH. Concentrated HCl was added to the reaction mixture
until the pH value was lowered to 2. The white precipitate was
filtered off and recrystallized from 2-propanol to yield 10 g of N-
(2-hydroxybenzyl)-NЈ-benzylethylenediamine diacetic acid hydro-
1
chloride as a white powder (m.p. 190 °C, yield 46%). Ϫ H NMR
([D6]DMSO): δ ϭ 3.14 (t, 2 H), 3.31 (t, 2 H), 3.42 (s, 2 H), 3.90 (2
ϫ s, 4 H), 4.27 (s, 2 H), 6.85 (td, 1 H), 6.95 (d, 1 H), 7.23 (m, 7
H). Ϫ C20H25ClN2O5 (ϩ 0.05 NaCl) (411.8): calcd. C 58.33, H
6.08, Cl 9.06, N 6.80, O 19.45; found C 58.21, H 6.09, N 6.53, O
19.86, Cl 9.20.
Potentiometric Titrations: All the measurements were made at 25
°C and the solutions were prepared with deionized water that had
been distilled twice. The ionic strength was fixed at I ϭ 0.1  with
sodium perchlorate (PROLABO puriss). The potentiometric titra-
tions were performed using an automatic titrator system, DMS 716
Titrino (Metrohm), equipped with a combined glass electrode (Me-
trohm, filled with saturated NaCl solution) and connected to an
IBM Aptiva microcomputer. The electrodes were calibrated to read
p[H] according to the classical method[31] (from titration of 0.1 
HClO4 by 0.1  NaOH). The ligand and its iron(III) complex of ca.
0.001  were titrated with standardized 0.025  sodium hydroxide.
Argon was bubbled through the solutions to exclude CO2 and O2.
Sodium hydroxide was prepared from 0.1  NaOH (Prolabo) and
was standardized. The titration data (120 points collected over the
pH range 2.5Ϫ10.5 for the ligand solution and 100 points collected
over the pH range 2.5Ϫ10 for the FeIIIϪligand solution) were re-
fined by the nonlinear least-squares refinement program SU-
PERQUAD[10] to determine the deprotonation constants (σfit in the
range 3Ϫ4). The pKan values were calculated from the cumulative
constants determined with the above program. The uncertainties in
the pKan values correspond to the standard deviations (1σ) in the
cumulative constants.
Experimental Section
Materials and Equipment: Ligands L1 and L2 were prepared ac-
cording to the procedure described previously.[8] All other com-
pounds were of reagent grade and were used without further puri-
fication. FeIII stock solutions were prepared by dissolving appropri-
ate amounts of ferric perchlorate hydrate (Aldrich) in standardized
HClO4 and NaClO4 solutions. The solutions were calibrated spec-
trophotometrically for ferric ions by using a molar extinction coef-
ficient of 4160 Ϫ1cmϪ1 at 240 nm.[30] Mass spectra were obtained
by HPLC/MS with a Fisons Platform mass spectrometer equipped
with an atmospheric pressure ion source in the electrospray ioniza-
Spectrophotometric Experiments: UV/Vis absorption spectra were
recorded with a PerkinϪElmer Lambda 2 spectrometer using
1.000-cm path-length quartz cells and connected to a microcom-
puter; the acquisition was made with UV Winlab software
(PerkinϪElmer). The temperature was maintained at 25 °C with a
PerkinϪElmer PTP-1 variable temperature unit. The ferric complex
with L3 was studied by spectrophotometry. The UV/Vis spectrum
of a solution containing equal amounts of ligand and FeIII (10Ϫ4
tion (ESI) mode. 1H NMR spectra were recorded with a Bruker ) was recorded as a function of pH over the range 1Ϫ8 (adjusted
500 MHz spectrometer.
478
with HClO4 or NaOH). An aliquot was taken from the solution
Eur. J. Inorg. Chem. 2001, 471Ϫ479

Iron(III) Complexation by New Aminocarboxylate Chelators
FULL PAPER
after each adjustment of the pH value and its spectrum was re-
corded. The pH measurements were made with a 713 Metrohm
digital pH meter equipped with a microelectrode. The ionic
strength was fixed at I ϭ 0.1  with NaClO4/HClO4. The spectro-
photometric data were analyzed using the LETAGROP-SPEFO
program.[12,13] The program uses a nonlinear least-squares method
to calculate the thermodynamic constants of the absorbing species
and their corresponding electronic spectra. The calculations were
performed using absorbance values from about 6Ϫ8 wavelengths
(between 400 and 600 nm). The range of values for the residual-
spectrophotometric and kinetic measurements with ligand L3. We
are grateful to Dr. Alain Deronzier (LEOPR, Grenoble, France) for
´
the use of the laboratory facilities, to Professor C. Beguin (LEDSS,
´
Grenoble, France), Professor M. Fontecave and Dr. S. Menage
(CBCRB, CEA, Grenoble, France) for their interest in these stud-
ies.
[1]
R. C. Hider, S. Singh, Proc. R. Soc. Edimburgh, Sect. B 1992,
99, 137Ϫ168.
[2]
R. C. Crichton, In Inorganic Biochemistry of Iron Metabolism,
Ellis Horwood Ltd., New York, 1991, p. 173.
squares sum [Σ(Aexp Ϫ Acalc)2] of the fits was 10Ϫ2Ϫ10Ϫ3
.
[3]
C. Herschko, Mol. Aspects Med. 1992, 13, 114.
[4]
J. B. Porter, E. R. Huens, R. C. Hider, Bailliere’s Clin. Haema-
Electrochemical Measurements: Electrochemical experiments were
carried out using a PAR model 273 potentiostat equipped with
a KippϪZonen x-y recorder. All experiments were run at room
temperature under argon in a glove-box. A standard three-electrode
cell was used. Potentials are referenced to Ag/AgCl, 3  NaCl,
aqueous reference electrode (ϩ0.21 V vs. NHE). A glassy carbon
disc electrode, used as a working electrode (5 mm diameter), was
polished with 1 µm diamond paste. The electrochemical behavior
of the FeIIIϪligand complex was studied by cyclic voltammetry
(CV) in an aqueous solution containing 0.1  NaClO4 as a sup-
porting electrolyte and buffered with tris-buffer for the experiments
performed at pH ϭ 7. For the experiments performed at pH ϶ 7,
the pH value was adjusted following controlled additions of con-
centrated NaOH or HClO4 aqueous stock solutions. The
FeIIIϪligand solutions were prepared by dissolving stoichiometric
amounts of ferric perchlorate and ligand in the electrolytic solu-
tions.
tol. 1989, 2, 257.
[5]
B. Halliwell, J. M. C. Gutteridge, Free Radicals in Biology and
Medicine, 2nd ed., Clarendon Press, Oxford, 1989.
[6]
J.-B. Galey, Adv. Pharmacol. 1996, 38, 167.
[7]
´
´
J.-B. Galey, J. Dumats, S. Genard, O. Destree, P. Pichaud, P.
Catroux, L. Marrot, I. Beck, B. Fernandez, G. Barre, M. Seite,
G. Hussler, M. Hocquaux, Biochem. Pharmacol. 1996, 51, 103.
[8]
´
´
S. Menage, J.-B. Galey, J. Dumats, G. Hussler, M. Seite, I.
Gautier-Luneau, G. Chottard, M. Fontecave, J. Am. Chem.
Soc. 1998, 120, 13370.
[9]
´
´
J.-B. Galey, O. Destree, J. Dumats, S. Genard, P. Tachon, J.
Med. Chem. 2000, 43, 1418.
[10]
P. Gans, A. Sabatini, A. Vacca, J. Chem. Soc., Dalton Trans.
1985, 1195.
[11]
[12]
[13]
W. R. Harris, A. E. Martell, Inorg. Chem. 1976, 14, 713.
L. G. Sillen, B. Warsquist, Ark. Kemi 1969, 31, 377.
M. Meloun, J. Havel, E. Hogfeldt, Computation of Solution
Equilibria, Ellis Horwood, Chichester, 1988.
W. R. Harris, A. E. Martell, Inorg. Chem. 1975, 14, 974.
R. J. Motekaitis, Y. Sun, A. E. Martell, M. J. Welch, Inorg.
Chem. 1991, 30, 2737.
[14]
[15]
Kinetics Studies: Kinetic measurements were performed with a
KINSPEC UV (BIO-LOGIC Company, Claix, France) stopped-
flow spectrophotometer equipped with a diode array detector (J &
M) and connected to a TANDON microcomputer. The kinetic data
were treated online with the commercial BIO-KINE program
(BIO-LOGIC Company, Claix, France). The ionic strength was
fixed at I ϭ 2  (NaClO4, HClO4) owing to the Hϩ concentrations
up to 1  and allowing comparisons with literature data. Ϫ Forma-
tion kinetic studies were carried out under pseudo-first-order con-
ditions at 25 °C with FeIII in excess with respect to the ligand. The
FeIII concentration spanned the range 5 ϫ 10Ϫ3Ϫ3 ϫ 10Ϫ2  for
each Hϩ concentration, which was over the range 0.05Ϫ0.15 . A
solution containing FeIII and Hϩ and a solution containing the
ligand L3 (2 ϫ 10Ϫ4 ) at the same ionic strength were mixed on
the stopped-flow apparatus. In each case, first-order kinetics were
observed. The reported rate constants are the average of about 8
replicate determinations (standard deviation in the range 1Ϫ2%).
Ϫ The acid hydrolysis kinetics of the FeIIIϪL3 complex was studied
under pseudo-first-order conditions in the presence of excess pro-
tons ([Hϩ] range 0.02Ϫ1.0 ) at 25 °C. The initial pH value of the
FeIIIϪL3 solution (7 ϫ 10Ϫ4 ) was 4.
[16]
[17]
[18]
R. M. Milburn, W. C. Vosburgh, J. Am. Chem. Soc. 1955, 77,
1352.
The term arising from the reverse reaction is negligible under
the experimental conditions of the hydrolysis reaction.
M. Birus, N. Kujundzic, M. Pribanic, Prog. React. Kinet. 1993,
18, 171.
M. Eigen, R. G. Wilkins, Adv. Chem. Ser. 1965, 55.
R. M. Fuoss, J. Am. Chem. Soc. 1958, 80, 5059.
E. Mentasti, E. Pelizetti, G. Saini, J. Chem. Soc., Dalton Trans.
1974, 1944.
[19]
[20]
[21]
[22]
[23]
Z. Zhang, R. B. Jordan, Inorg. Chem. 1996, 35, 1571.
M. Birus, Z. Bradlic, N. Kujundzic, M. Pribanic, P. C. Wilkins,
R. G. Wilkins, Inorg. Chem. 1985, 24, 3980.
M. T. Caudle, A. L. Crumbliss, Inorg. Chem. 1994, 33, 4077.
[24]
[25]
´
F. Thomas, C. Beguin, J.-L. Pierre, G. Serratrice, Inorg. Chim.
Acta 1999, 291, 148.
[26]
[27]
[28]
[29]
A.-M. Albrecht-Gary, S. Blanc, N. Rochel, A. Z. Ocaktan, M.
A. Abdallah, Inorg. Chem. 1994, 33, 6391.
B. Monzyk, A. L. Crumbliss, J. Am. Chem. Soc. 1982, 104,
4921.
M. Birus, Z. Bradlic, G. Krznaric, N. Kujundzic, M. Pribanic,
P. C. Wilkins, R. G. Wilkins, Inorg. Chem. 1987, 26, 1000.
´
´
J.-B. Galey, O. Destree, J. Dumats, P. Pichaud, J. Marche, S.
´
Genard, G. Bracciolli, L. Le Capitaine, H. Plessix, L. Bram-
billa, O. Cantoni, Free Rad. Biol. Med. 1998, 25, 881.
R. Bastian, R. Weberling, F. Palilla, Anal. Chem. 1956, 28, 459.
A. E. Martell, R. J. Motekaitis, Determination and Use of
Stability Constants, 2nd ed., VCH Publishers, New York, 1992.
Received May 3, 2000
[30]
[31]
Acknowledgments
´
The authors thank Celine Carresi for performing the potentiomet-
ric measurements and Arnaud Pierra for performing the UV/Vis
[I00179]
Eur. J. Inorg. Chem. 2001, 471Ϫ479
479