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Plant Physiol. (1999) 119: 1107-1114
Nicotianamine Chelates Both FeIII and
FeII. Implications for Metal Transport in
Plants1
Nicolaus von Wirén2,
Sukhbinder Klair,
Suhkibar Bansal,
Jean-Francois Briat,
Hicham Khodr,
Takayuki Shioiri,
Roger A. Leigh*, and
Robert C. Hider
Department of Pharmacy, King's College London, Manresa Road,
London SW3 6LX, United Kingdom (N.v.W., S.K., S.B., H.K., R.C.H.); Laboratoire de Biochimie et Physiologie Moléculaire des Plantes,
Institut National de la Recherche Agronomique, Place Viala, F-34060
Montpellier, France (J.-F.B.); Faculty of Pharmaceutical Sciences,
Nagoya City University, Tanabe-dori, Mizuho-ku, Nagoya 467, Japan
(T.S.); and Biochemistry and Physiology Department, IACR-Rothamsted,
Harpenden, Hertfordshire AL5 2JQ, United Kingdom (R.A.L.)
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ABSTRACT |
Nicotianamine
(NA) occurs in all plants and chelates metal cations, including
FeII, but reportedly not FeIII. However, a
comparison of the FeII and ZnII affinity
constants of NA and various FeIII-chelating
aminocarboxylates suggested that NA should chelate FeIII.
High-voltage electrophoresis of the FeNA complex formed in the presence
of FeIII showed that the complex had a net charge of 0, consistent with the hexadentate chelation of FeIII.
Measurement of the affinity constant for FeIII yielded a
value of 1020.6, which is greater than that for the
association of NA with FeII (1012.8). However,
capillary electrophoresis showed that in the presence of
FeII and FeIII, NA preferentially chelates
FeII, indicating that the FeIINA complex is
kinetically stable under aerobic conditions. Furthermore, Fe complexes
of NA are relatively poor Fenton reagents, as measured by their ability
to mediate H2O2-dependent oxidation of
deoxyribose. This suggests that NA will have an important role in
scavenging Fe and protecting the cell from oxidative damage. The pH
dependence of metal ion chelation by NA and a typical phytosiderophore,
2 -deoxymugineic acid, indicated that although both have the ability to
chelate Fe, when both are present, 2-deoxymugineic acid dominates the chelation process at acidic pH values, whereas NA dominates at alkaline
pH values. The consequences for the role of NA in the long-distance
transport of metals in the xylem and phloem are discussed.
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INTRODUCTION |
NA,
2(S),3(S),3"(S)-N-[N-(3-amino-3-carboxypropyl)-3-amino-3-carboxypropyl]-azetidine-2-carboxylic
acid (Fig. 1, structure I), and MA,
2(S),2(S),3(S),3"(S)-N-[3-carboxy-(3-carboxy-3-hydroxypropylamino)-2-hydroxypropyl]-azetidine-2-carboxylic acid (structure II a) are two structurally similar molecules
synthesized by plants. Although both compounds have roles in the
acquisition and transport of Fe, their species distribution and
physiological functions are distinct. MA and some closely related
compounds, termed PS, are made only by graminaceous monocotyledonous
plants and are secreted from the roots of Fe-deficient grasses to
mobilize Fe from insoluble sources (Takagi, 1976 ;
Sugiura and Nomoto, 1984). These compounds chelate
FeIII (Sugiura and Nomoto, 1984), and the
FeIIIPS complex is probably the form in which Fe
is taken up by the roots of grasses (Römheld and Marschner, 1986;
von Wirén et al., 1995).
By contrast, NA is made by all plants and is present in various plant
organs (Stephan et al., 1994; Walter et al., 1995) but is not secreted.
Instead, it is thought to have a role in the internal transport of Fe
and other metals (Stephan et al., 1994, 1996; Pich and Scholz, 1996).
Evidence in support of this role is that the concentrations of NA in
the phloem correlate with those of Fe and other metals, and the
NA-synthesis-defective tomato mutant chloronerva has a
phenotype indicative of Fe deficiency (Pich and Scholz, 1996; Stephan
et al., 1996). Despite its structural similarity to MA and other PS, NA
is proposed to fulfill its physiological role in Fe trafficking by
chelating FeII and not
FeIII (Stephan and Scholz, 1993). Although
FeIIINA has been demonstrated by
electron-spin-resonance spectroscopy (Sugiura and Nomoto, 1984), K for
the formation of this complex has not been successfully measured and so
it has been assumed to be physiologically unimportant. In contrast, K
of NA for FeII has been successfully measured
(Bene et al., 1983; Anderegg and Ripperger, 1989). Thus, it has
been concluded that the FeII complex is the
only significant one in biological systems, despite the
electron-spin-resonance spectroscopy evidence for the formation of an
FeIIINA complex.
It is important to establish whether the FeIIINA
complex exists under physiological conditions because it could
fundamentally affect the conclusions about the forms in which Fe and
other metal ions are transported within plants. If NA functions only as
an FeII chelator, Fe taken up by grasses as
FeIIIPS complexes will need to be reduced before
formation of the FeIINA complex for subsequent
internal transport to other parts of the plant. However, if NA chelates
FeIII, then it may be possible for the Fe to be
donated directly to it by the FeIIIPS complex. In
addition, if NA binds FeIII more tightly than
FeII and other divalent metal ions, this will
affect the relative abundance of different NA-metal complexes, possibly
altering conclusions about the role of NA in the internal trafficking
of these metals (Stephan et al., 1996). Here we provide evidence that
NA does indeed form an FeIII complex under
physiological conditions, but we also show that the
FeIINA complex has an unexpected kinetic
stability. The role of NA in protecting cells from oxidative damage
through scavenging of Fe and in metal trafficking are considered in
relation to these properties.
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MATERIALS AND METHODS |
Chemicals
NA (structure I), MA (structure II a), and DMA (structure II b)
were chemically synthesized as previously described and were at least
95% pure as judged by 1H-NMR (Shioiri et al.,
1995). EDTA (structure III a, >99%), HEDTA (structure III b, >98%),
and DTPA (structure IV, >98%) were purchased from Sigma. DFO
(structure V, >90%) was a gift from Novartis (Basel, Switzerland). All other chemicals were from Aldrich. For the
preparation of labeled chelates an aliquot of 10 mM
Fe(NO3)3 containing 20 kBq
59FeCl3 (Amersham) was
added to 10 µL of 12 mM chelator solution and adjusted to
pH 7.0 with 100 mM Mops-KOH.
High-Voltage Electrophoresis
A sheet of filter paper (grade 1F, Munktell, Stockholm,
Sweden) was placed on a cooling plate covered by acetate foil in an electrophoresis cuvette (Multiphor II, Pharmacia). The filter paper was
positioned so that both ends dipped in buffer solution (0.1 M Mops-KOH, pH 7.0) and the paper was pre-run at 400 V for 20 min. Small pieces of filter paper (10 × 4 mm) were loaded with approximately 10 µL of 1 mM
59Fe-chelate solution (2 kBq) and were then
placed on a start line in the middle of the paper sheet. Separation was
achieved in the dark at a constant 400 V at 10°C. After 60 min,
electrophoresis was stopped, and the paper sheet was immediately dried
with a hair drier and cut into 8-mm-wide strips parallel to the start line. The amount of radioactivity on the paper strips and on the deposit paper was determined in a  -counter (LKB, Bromma,
Sweden) by dry Cerenkov counting. For a comparison of migration
distances, results were expressed in counts per minute per strip. The
net charge was calculated from the migration distances using the
59FeIII and
65ZnII complexes of EDTA,
HEDTA, and DTPA as the standards.
pKa Determination
Equilibrium constants of protonated ligands were determined using
an automated, computerized system capable of simultaneously analyzing
spectrophotometric and potentiometric measurements. A blank titration
of 0.1 M KCl was performed to determine the electrode zero
using Gran's plot method (Gran, 1952). A combined pH electrode
(Sirius Analytical Instruments Ltd., Forest Row, East Sussex, UK) was
used to calibrate the electrode zero. The solution (0.1 M
KCl, 25 mL), contained in a jacketed titration cell to maintain the
temperature at 25°C ± 0.5°C, was under an argon atmosphere and was
acidified with 0.15 mL of 0.2 M HCl. Titrations were
carried out against 0.2 M KOH added in 0.01-mL increments
dispensed from a Metrohm 665 dosimat (Metrohm Ltd., Buckingham, UK).
The titration was repeated in the presence of ligand. The data obtained
from the titrations were analyzed by the TITRFIT program, a modified
version of NONLIN (Taylor et al., 1988).
Determination of K
The K for the FeIII-ligand interaction was
determined by a spectrophotometric competition study of
ligand-FeIII-maltol using the automated
system described above. The FeIII complexes of
the ligand were prepared in a 10:1 ligand:Fe molar ratio (total
Fe concentration = 4.4 × 10 5 M)
in 0.1 M Mops-KOH buffer, pH 7.4. This solution was
then titrated against maltol (3-hydroxy-2-methylpyran-4-one), resulting
in the dissociation of FeIIINA and the formation
of the orange FeIIImaltol complex. The resulting
spectrophotometric data were inserted into the COMPT1 program (H. Khodr, unpublished data) to evaluate K of the complex. Speciation and
pM plots were obtained by running the program SPECIAZ1 (H. Khodr,
unpublished data). These programs require concentrations of metal and
ligand, K(FeIII) values of complexes, K values
for FeIII-OH interactions, and
pKa values of the ligand. Between pH 1.0 and
9.0, the dominant FeIII-OH species are
[FeIII(OH)]2+,
[FeIII(OH)2]+,
[FeIII(OH)3], and
[FeIII(OH)4].
CE
Separations were performed at 25°C in fused silica capillaries
(37 cm long x 75 µm i.d.; Metal Services Ltd., Worcester, UK) on a
PACE 5510 apparatus (Beckman) equipped with a photodiode array detector
and using Nouveau Gold software (Beckman) for data acquisition. The
capillary was rinsed with 0.1 mM NaOH and then with water
for 5 min before equilibration for 20 min with the carrier electrolyte.
These procedures were repeated but for only 2 min each between each
separation. The solutes were injected in the hydrodynamic mode by
overpressure (3.45 kPa). Electroosmotic flow was evaluated from the
migration time of formamide. For the separation of NA-metal complexes,
the concentration of NA was 2.5 mM and that of the metal
ion was 3 mM. Borate buffer (50 mM, pH 9.2) was
the carrier electrolyte and separation was for 12 s at 667 V
cm1. For the competition studies between NA and
citrate, the conditions were the same, except 25 mM Na
citrate was added and the carrier electrolyte was 50 mM
phosphate buffer, pH 5.5.
Measurement of Fenton Activity of FeNA complexes
Fenton activity was determined in the presence of 1.5 mM H2O2 and 2.8 mM deoxyribose by incubation of various FeNA complexes (total Fe concentration = 20 µM, total NA
concentration = 100 µM) in 50 mM
phosphate buffer, pH 7.4, at 37°C. (Halliwell et al., 1987). The
formation of thiobarbituric acid-reactive substances from deoxyribose
was quantified by measurement of A532
(Halliwell and Gutteridge, 1981). Results are the mean of five
determinations and are corrected for absorbance in the absence of
deoxyribose. When present, ascorbic acid was used at 100 µM.
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RESULTS AND DISCUSSION |
Both NA and MA have six ligands for Fe complexation, and the
distances between the groups facilitate octahedral coordination of Fe
and the formation of three 5-membered and two 6-membered chelate rings
(Ripperger and Schreiber, 1982). The only differences between
NA and MA are the hydroxyl/amino substitution on C3" and the presence
of a hydroxyl group on C2 of MA (Fig. 1). Two observations led us to
query the conclusion that there is no interaction of NA with
FeIII.
The first was a comparison of NA's K for FeII
and ZnII (Anderegg and Ripperger, 1989) with
those of a closely related group of oligodentate aminocarboxylate
ligands, including MA and DMA, which are known to chelate
FeIII. The ratios of the affinities of these
ligands for ZnII and FeII
are similar to that for NA and all fall into a tight range, 1.12 to
1.26 (Table I). The value of this ratio
for NA (1.20) is close to the mean for the group (1.19). Similarly, the
ratios of affinities for FeIII and
FeII fall between 1.62 and 1.77. Using these
ratios to predict the K for the interaction between
FeIII and NA suggests a value that falls in the
range of 1020 to 1022. This
estimated value is much higher than that reported for the FeII-NA interaction
(1012.1-1012.8;
Bene et al., 1983; Anderegg and Ripperger, 1989) or for the FeIII-MA interaction
(1018.2; Sugiura and Nomoto, 1984). Furthermore,
molecular modeling of a putative FeIIINA complex
and comparison with that for FeIIIDMA showed that
the two structures are very similar, with no obvious strain in the
FeIIINA complex (Fig.
2).
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| Figure 2.
Computer-generated models of the FeIII
complexes of NA and DMA. The coordinates are based on the
CoIII complex of MA (Sugiura et al., 1981) and the
ionization state demonstrated by the high-voltage electrophoresis data
in Figure 3.
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Second, using high-voltage electrophoresis of NA and MA in the presence
of 59FeIII at pH 7.0 under
aerobic conditions in the dark (so FeII would not
be present), we measured net charges of 0 and  1.0 for the FeNA and
FeMA complexes, respectively (Fig. 3).
These values are entirely consistent with a hexadentate structure for both molecules chelating FeIII (Fig. 2). The lack
of charge on the NA complex is due to the compensation of the negative
charges of the carboxylate groups by the FeIII.
In the case of MA, the extra negative charge is associated with the
deprotonation of the terminal OH function in the presence of
FeIII (Sugiura et al., 1981). These net charges
specifically exclude pentadentate (terminal amino function not involved
in coordination) and tetradentate ( -amino acid functions not
involved in coordination) structures, which for NA would both have net
charges of +1. The FeIII complex of DMA also
possesses a net charge of 1.0 at pH 7.0 (data not shown).
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| Figure 3.
Net charge of FeIII complexes of NA
( ) and MA ( ) as determined by high-voltage paper electrophoresis
at pH 7.0 under aerobic conditions in the dark.
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In view of these strong indications that NA can bind
FeIII, a spectrophotometric titration procedure
was used to determine directly both the pKa
values and the K for FeIIINA by measuring the
ability of NA to compete with maltol for FeIII.
Titration data for NA are shown in Figure
4. Maltol is a powerful FeIII chelator that generates an orangecolor on
complexation with FeIII (Hider and Hall, 1991).
This method avoids the difficulty of Fe(OH)3
precipitation reported by Bene et al. (1983). To verify the
accuracy of the method, pKa and
FeIII K were measured for DMA and EDTA. The
measured values agreed well with those from the literature (Table
II). The method gave a value of
1020.6 for the K for the interaction of NA with
FeIII (Table II). This is in agreement with the
predictions from the data from Table I, providing a ratio for the K for
FeIII and FeII of 1.61, and
indicates that, contrary to current dogma, NA binds FeIII with an affinity 108
times greater than that with which it binds FeII.
This finding also offers an explanation for the reported redox potentials of the Fe complexes of MA and NA (Table
III; Sugiura and Nomoto, 1984). More
negative redox potentials indicate selectivity for
FeIII over FeII as
exemplified by DFO, a highly specific
FeIII-chelating bacterial trishydroxamate
siderophore (KFeIII = 30.6, KFeII = 7.2; Anderegg et al.,
1963). The redox potential of DFO is significantly more negative than
that of NA or MA, consistent with DFO's relative inability to chelate
FeII.
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| Figure 4.
Spectrophotometric titration against maltol of a
43.5 µM solution of FeIIINA complex.
FeIIINA dissociates with increasing concentrations of
maltol, leading to the formation of the orange FeIIImaltol
complex. The concentrations of maltol used are indicated.
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Table II.
Measured and published pKa values and K
for FeIII complexes of NA, DMA, and EDTA
Published values (where available) are in parentheses: Smith and
Martell (1989) (DMA), Murakami et al. (1989) (NA), and Anderegg and
Ripperger (1989) (EDTA).
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The efficiency of chelation by different ligands is best compared using
pM values, which provide a measure of the free metal ion concentration
in the presence of ligand (Hider, 1995). The higher the pM value, the
more powerful the chelator. The influence of pH on the
pFe3+ values for NA, DMA, EDTA, and DFO is
presented in Figure 5A. It is clear that
NA and DMA are both markedly weaker chelators of
FeIII than either EDTA or DFO. They are just able
to compete with the hydroxide anion over the pH range of 3.0 to 8.0. In
contrast, although the K values of the FeII
complexes of NA and DMA are much lower than those of the corresponding FeIII complexes (Tables I and II), the
pFe2+ values indicate that NA is capable of
efficiently competing with hydroxide at pH values greater than 6.5 (Fig. 5B). DMA is an inferior ligand for FeII,
only beginning to compete effectively with the hydroxide anion at pH
values greater than 7.0. The plot also indicates the relative inability
of DFO to complex FeII. In summary, in aqueous
solution NA binds FeIII over the pH range of 4.0 to 9.0 and binds FeII over the pH range of 6.0 to
9.0, whereas in general, DMA and PS bind FeIII
over the pH range of 2.5 to 8.5 and bind FeII
over the pH range of 7.5 to 9.0.
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| Figure 5.
The influence of pH on pFe3+ (A) and
pFe2+ (B) values for hydroxide, NA, DMA, EDTA, and DFO. The
total Fe concentration was 1 µM and the total ligand
concentration was 10 µM. The higher the pM value, the
more effective the ligand.
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Although EDTA forms a complex with FeII, this
rapidly autoxidizes in air and thus the FeIII
complex always dominates under aerobic conditions (Zang and van Eldik,
1990). By virtue of the high affinity of NA for
FeIII, the FeIINA complex
would be expected to behave similarly, which would mean that the
complex would be unlikely to have a transport function in aerobic plant
tissues. Therefore, the stability of the FeIINA
complex was investigated using CE. Unlike the FeII
complexes of EDTA and DTPA, the FeIINA complex
was found to be quite stable, possessing a running time similar to that
of CuIINA and quite distinct from that of the
FeIIINA complex (Fig.
6). There was no evidence of autoxidation
of the FeIINA complex, even when
O2 was bubbled through the solution. This is
surprising in view of the values of the stability constants of the
FeII and FeIII complexes
(see above) and indicates that, once formed, the
FeIINA complex is kinetically stable. To test
this, equimolar amounts of FeII and
FeIII were added to NA at pH 7.0 (final ratio of
total Fe:NA was 2:1) and the complexes separated by CE. Only the
FeII complex was detected (not shown). Thus, at
pH 7.0, NA will preferentially scavenge FeII even
in the presence of FeIII, and the
FeIINA complex will persist by virtue of its
kinetic stability. Consistent with this finding, measurements of the
potency of FeNA complexes as Fenton reagents, determined from their
ability to mediate
H2O2-dependent oxidation of
deoxyribose, showed that they are not highly active Fenton reagents
(Table IV) and do not redox cycle
efficiently. These physicochemical properties will allow NA to scavenge
FeII, preventing
FeII-mediated production of hydroxyl radicals
from H2O2 via the Fenton reaction (Guerinot and
Yi, 1994; Henle and Linn, 1997) and protecting the cell from Fe-induced
oxidative damage.
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| Figure 6.
Migration times of NA and its Fe complexes as
measured by CE. A, NA alone; B, FeIINA; C,
FeIIINA.
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Table IV.
Fenton activities of Fe complexes of NA in
comparison with those of Fe-EDTA and Fe-phosphate
Values are means ± SE of five separate
determinations.
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Physiological Implications
To determine the possible physiological consequences of these
observations for metal trafficking in plants, we used the
pKa and measured K of NA to calculate the
concentrations of different complexes under different conditions and
compared these with other physiologically important
FeIII chelators, such as PS and citrate. First,
we examined the effect of pH on the relative abilities of NA and DMA to
chelate FeIII to determine whether the
hydroxyl/amino substitution at C3" has any major significance that
could explain the different physiological roles that these molecules
play in Fe acquisition and trafficking (see the introduction). Although
both compounds chelate FeIII over a broad pH
range, FeIIINA is more acid labile than DMA and,
conversely, is more stable at alkaline pH values (Fig.
7, A and B). As a result, in an
equilibrium mixture FeIIIDMA will dominate at pH
3.0 to 6.0 and FeIIINA at pH 7.0 to 9.0 (Fig.
7C); this agrees with the data presented in Figure 5A. The
FeIINA complex is also acid labile (Fig. 7D;
Stephan et al., 1996). These observations are consistent with the
proposed physiological roles of NA and PS. In graminaceous plants PS
need to be able to chelate FeIII over a wide pH
range, from alkaline values in calcareous soils where Fe deficiency is
most common (Guerinot and Yi, 1994) to slightly acidic values in the
root apoplast where the FeIIIPS complex is
transported across the root plasma membrane. This behavior is supported
by the observation that PS can mobilize Fe and mediate its uptake over
this pH range (Römheld and Marschner, 1986; Treeby et al., 1989).
Once the FeIIIPS complex has been absorbed, the
greater relative stability of the NA complex at cytosolic pH values
(7.2-7.5) means that NA will more efficiently chelate Fe in this
compartment. Any metabolism or reduction of the
FeIIIPS complex that occurs to enhance Fe release
will facilitate the rate of transfer, and the ability of NA to chelate
both FeII and FeIII will
result in Fe being scavenged rapidly irrespective of which form becomes
available.
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| Figure 7.
Computer simulations of the pH dependence of the
Fe complexes of NA and DMA. A, FeIII and NA; B,
FeIII and DMA; C, competition for FeIII between
NA and DMA; D, FeII and NA. In all cases, the total Fe
concentration is 1 µM, and NA or DMA is present at 10 µM.
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PS also facilitate both the mobilization of Zn2+
and Cu2+ (Treeby et al., 1989) and the uptake of
Zn2+ (von Wirén et al., 1996), and NA is
proposed to traffic these metals within the plant (Stephan et al.,
1996). Thus, considerations similar to those mentioned above occur in
relation to the relative stability of the PS and NA complexes of these
metals. As with Fe, the NA complexes with ZnII
and CuII are acid labile, with the
CuII complex being stable to slightly lower pH
values than the ZnII complex (Fig.
8, A and B; Stephan et al., 1996).
However, when DMA is also present, the equilibrium metal distribution
is different for the two metals. For ZnII, the NA
complex dominates virtually over the entire physiologically relevant pH
range (Fig. 8C), whereas the DMA complex dominates for
CuII (Fig. 8D).
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| Figure 8.
Computer simulations of the pH dependence of the
Zn and Cu complexes of NA and MA. A, ZnII and NA; B,
CuII and NA; C, competition for ZnII between NA
and DMA; D, competition for CuII between NA and DMA. The
total concentration of ZnII or CuII is 1 µM, and NA or DMA is present at 10 µM.
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The major role that has been proposed for NA is as a facilitator of the
long-distance transport of a number of metals in the phloem (Stephan
and Scholz, 1993; Schmidke and Stephan, 1995; Stephan et al., 1996) and
of CuII in the xylem (Pich et al., 1994; Pich and
Scholz, 1996). Some of these conclusions have been based on the
assumption that NA chelates FeII but not
FeIII, and so it is important to consider whether
these general conclusions would change with the knowledge that NA binds
FeIII with high affinity. Therefore, we
calculated the equilibrium concentrations of different metal complexes
in conditions approximating those in the xylem and phloem. Calculations
for the xylem (NA = 20 µM, citrate = 150 µM, Fe = 40 µM, Zn = 5 µM, and Cu = 5 µM, pH 5.5; Pich and
Scholz, 1996; A. Pich, personal communication) showed that all Fe is
complexed with citrate. This is expected on the basis of the speciation
plots, which show that for both FeIII and
FeII, citrate complexes totally dominate at pH
5.5 (Fig. 9). However, unlike
FeIINA, the FeIIcitrate
complex rapidly autoxidizes (Harris and Aisen, 1973). We have confirmed
with CE that citrate removes Fe from FeIINA at pH
5.5 (Fig. 10). Under these conditions
Fe is removed from NA with an estimated half-life of less than 5 min
and thus must be converted to FeIIIcitrate;
hence, this complex will dominate even if FeII is
the major form in which Fe is loaded into the xylem. As there are
reports of citrate concentrations much higher than 150 µM in the xylem (e.g. up to 1700 µM; Tiffin, 1996; White et
al., 1981), the speciation plots were also calculated with a citrate concentration of 1500 µM. This made no difference to the
conclusions about the complexation of Fe at pH 5.5 but slightly
increased the preponderance of Fe citrate complexes at alkaline pH
values (not shown). In contrast to Fe, NA was found to be more
important than citrate in the chelation of both
ZnII and CuII. With the
metal and ligand concentrations given above and with citrate at 150 µM, the calculations showed that NA chelates 50% of the
ZnII and all of the CuII.
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| Figure 9.
Computer simulations of the competition between
citrate and NA for FeIII (A) and FeII (B).
Metal and ligand concentrations are those estimated for the xylem:
total Fe = 40 µM, citrate (Cit) = 150 µM, and NA = 20 µM.
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| Figure 10.
Competition for FeII between NA and
citrate at pH 5.5, as monitored by CE. Citrate was added at 0 min and
the amount of FeIINA was measured by CE at the times
indicated. The inset shows traces from the CE at 2 min (A) and 12 min
(B). In each trace, the left peak is NA and the right peak is
FeIINA.
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Calculations for the phloem (NA = 130 µM,
citrate = 1500 µM, Fe = 45 µM,
Zn = 30 µM, and Cu = 15 µM, pH
8.0; Schmidke and Stephan, 1995; value for citrate from Jeschke et al.,
1986) indicated that all of the metal ions are present as their NA
complexes; citrate complexes are predicted to play no role in this
pathway. Whether FeIINA or
FeIIINA is the major Fe complex in the phloem is
unclear, but if Fe is loaded into the phloem as either
FeIINA or FeII, then the
kinetic stability of the FeIINA complex will
probably ensure that this complex dominates.
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CONCLUSIONS |
It is now clear that studies of the physiological role of NA in
plants have wrongly assumed that NA does not chelate
FeIII. The studies in this paper show that NA is
an effective chelator of FeIII, with an K
appreciably higher than that of the chemically related PS, which are
established FeIII chelators (Sugiura and
Nomoto, 1984). However, the FeIINA complex
possesses an unusual kinetic stability and does not rapidly autoxidize
to FeIIINA at physiological pH values. Clearly
this is an important property of the complex governing the evolution of
this molecule as an Fe chelator in plants; the significance of this
property requires further investigation. By demonstrating the existence
of the FeIIINA complex and analyzing the pH
dependence of metal-NA complexes, this study has provided a firmer
physicochemical basis for further physiological investigations of the
role of NA in chelation of Fe and other metals in plants. It has also
shown that scavenging of Fe by NA may be important in protecting the
cell from oxidative damage resulting from the Fenton reaction.
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FOOTNOTES |
1
The work was supported by a grant from the
Biotechnology and Biological Sciences Research Council (BBSRC) of the
United Kingdom and by a short-term fellowship to N.v.W. from the joint
BBSRC/Institut National de la Recherche Agronomique collaboration
scheme. IACR is grant-aided by the BBSRC.
2
Present address: Institut für Allgemeine
Botanik, Universität Tübingen, Morgenstelle 1, D-72076
Tübingen, Germany.
*
Corresponding author; e-mail RL225{at}cam.ac.uk; fax 44-1223-333953.
Copies of the computer program mentioned in the paper are available
from R.C.H. or H.K.
Received July 8, 1998;
accepted November 30, 1998.
| |
ABBREVIATIONS |
Abbreviations:
CE, capillary electrophoresis.
DFO, desferrioxamine.
DMA, deoxymugineic acid.
DTPA, diethylenetriaminepentaacetic acid.
HEDTA, N-hydroxyethylethylenediaminetriacetic acid.
K, affinity
constant(s).
MA, mugineic acid.
NA, nicotianamine.
pM, log10 free metal ion concentration.
PS, phytosiderophore(s).
| |
ACKNOWLEDGMENTS |
We thank Dr. P. Evans (King's College, London, UK) for making
the measurements of Fenton activities and Novartis for the supply of
DFO.
| |
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Abstract
Introduction
