|
|
|
|||||||
|
|||||||||
J Biol Chem, Vol. 274, Issue 19, 13375-13383, May 7, 1999
From the The recent use of calcein (CA) as a fluorescent
probe for cellular iron has been shown to reflect the nutritional
status of iron in mammalian cells (Breuer, W., Epsztejn, S., and
Cabantchik, Z. I. (1995) J. Biol. Chem. 270, 24209-24215). CA was claimed to be a chemosensor for iron(II), to
measure the labile iron pool and the concentration of cellular free
iron(II). We first study here the thermodynamic and kinetic properties
of iron binding by CA. Chelation of a first iron(III) involves one
aminodiacetic arm and a phenol. The overall stability constant log
Physiologists have described many examples of cells that present
symptoms of iron deficiency, although containing large amounts of iron
(1). In order to explain the capricious biological efficiency of
cellular iron, an ill defined key operator of iron metabolism, the so
called "chelatable" or "labile iron pool"
(LIP),1 has been invoked (2).
The LIP is the available cellular iron to which has been assigned at
least four functions: (a) cellular iron transport,
(b) expression of iron regulatory genes (transferrin receptor, ferritin), (c) control of the activity of iron
containing proteins, (d) catalysis of the Fenton reactions
(3). How much iron is available? How tightly is it bound and to which
biological ligands? Cabantchik et al. (4) have recently used
a fluorescent probe that may be an opportune tool to measure the LIP.
They have shown that calcein (CA; see structure formula in Fig. 3) is
dynamically sensitive to metabolic changes of iron pools, under various
nutritional conditions. Its use is therefore a methodological
breakthrough. Nevertheless, they assumed that thermodynamic data
concerning iron (III and II) complexation with calcein were the same as
those of EDTA, because calcein and EDTA both carry diaminetetraacetic chelating arms (4). By the use of acetoxymethyl esters of CA, they
could discriminate endocellular signals. Since calcein chelates cellular iron which is shared by all the complexes of the same compartment, the level of fluorescence of free calcein reflects the
variations of the available iron of the compartment (4-9). Calcein was
claimed to reflect the ferrous iron pool (10).
Such an approach is reminiscent of the model based on the notion of
metal buffers and pFe
calculations,2 by which
ligands and complexes tend to keep metal ions in homeostatic availability (11-17); biochemical iron exchanges can be regarded as
the result of competitions for iron between cellular ligands, according
to their various affinities. The biosynthesis of iron ligands during
growth of cells creates a driving force for the nutritional uptake of
the metal. Then, after iron uptake, the competition of iron complexes
with free ligands leaves only trace amounts of iron ions, in
equilibrium with all the iron potential ligands. Within each cellular
compartment, in which diffusion occurs freely, a single overall
concentration of unchelated residual iron results from these
multiple equilibria; it can be expressed as pFe. The term pFe is the
best quantitative expression of the overall concentrations of
unchelated iron in solution, in specific pH and ionic strength
conditions. CA is a fluorescent chemosensor that could reflect these equilibria.
First attempts to measure pFeIII in plant cells have
been made by competition experiments between radiolabeled standard iron buffers and cellular ligands (18), but these measurements were done
after disruption of cellular compartments and therefore had lost most
of their physiological significance. Nondestructive tools are needed
for pFeIII and pFeII determinations.
The aim of this work is to complement the cytological approach which
uses calcein as a chemosensor, by establishing the thermodynamic and
kinetic data concerning iron(III and II) complexation with calcein. We
then discuss the use of CA as a fluorescent probe for studying the
pools of ferrous and ferric iron. A new application using CA for
measurements of the kinetics of iron(III) nutrition of plant cells is
developed, and a physiological study of CA as a nutritional siderophore
is presented.
Potentiometric and UV-visible Spectrophotometric
Experiments--
All the measurements were made at 25 °C. The ionic
strength was fixed at 0.1 M with sodium perchlorate.
Potentiometric titration employed a DMS 716 Titrino (Metrohm) equipped
with glass and calomel electrodes and connected to an IBM Aptiva
microcomputer. The electrodes were calibrated to read pH according to
the classical method. A calcein solution and a 1:1 Fe(III):calcein
solution of ~0.001 M were titrated with standardized 0.05 M sodium hydroxide. The titration data were refined by the
nonlinear least squares refinement program SUPERQUAD (19) to determine
the equilibrium constants. UV-visible spectra were recorded on a Lambda
2 Perkin-Elmer spectrophotometer. Light path length is 1 cm. The
acquisition was made with the UV Winlab Perkin-Elmer software.
Temperature was maintained at 25 °C with the variable temperature
unit. For calcein-iron spectrophotometric titrations, 0.1 M
iron(III) perchlorate in 0.1 M HClO4 solution was added to a 10 Stopped Flow Experiments--
A Biologic SFM3 stopped flow
module (Claix, France), controlled by a Tandom computer using Biokine
software, is used for stopped flow kinetic experiments. Its temperature
control unit (M3 Lauda) is set at 25 °C.
EPR Experiments--
EPR studies have been conducted using a ESP
300E Bruker apparatus with a variable temperature unit. Spectra were
treated by using either the software of the apparatus or the winEPR
software (Bruker). Each sample (200 µl) contained 10% glycerol, 0.1 M Tris-Cl buffer, pH 7.2, 10 Electrochemical Studies--
The electrochemical investigations
by cyclic voltammetry (CV) were performed in 0.05 M
Tris-Cl, pH 7.2, 0.05 M NaClO4 aqueous solution
containing millimolar concentration of the iron(III)-calcein complex.
CV curves were recorded using a EGG 273 Potentiostat coupled with a
Kipp & Zonen x-y recorder. Electrochemical experiments were performed
in a three-compartment cell, at room temperature, under an argon
atmosphere. Potential are referred to an aqueous SCE reference
electrode. The working electrode was a vitrous carbon disc (5 mm
diameter) polished with 1-µm diamond paste.
Calcein + FeII Preparations--
The 1:1
FeII + CA solution for UV-visible spectrophotometry and
fluorimetric experiments has been prepared under argon in a glove box,
starting from the solids (CA, FeII as
(NH4)2(SO4)2Fe·6H2O)
and a degassed 0.1 M Tris-HCl, pH 7.2, solution. The
FeII + CA solution was diluted to the appropriate
concentration and placed into an hermetically closed quartz cell. Cells
were taken out to the fluorimeter or the spectrophotometer, and their
spectra were recorded. This prevents further Fe2+
oxidation, but not ferric contaminants from the powders.
Cell Culture and Growth Measurements--
Plant cells were grown
in 24-well sterile cell culture plates (Nunc); each independent axenic
culture contains 1.5 ml of culture medium. This medium (MS) was a
modified Murashige and Skoog (21): for 1000 ml of medium: 4.3 g of
macro- and micro-elements powder (provided by Duchefa, catalog number
M0221, Haarlem, Netherlands), 10 ml of H2KPO4
(20 g/liter), 0.5 ml; kinetin (0.1 g/l), 0.5 ml; (2,4-dichlorophenoxy)acetic acid (0.2 g/liter); 1 ml × 1000 vitamin solution (Duchefa M 0409), and 30 g of saccharose.
Iron-free MS medium (MS-Fe) has the same content, made from reagent
grade products, omitting FeIIIEDTA. Calcein is commercially
available (Molecular Probes, Eugene, OR). Wells were inoculated at an
absorbance of 0.003 (~
The measurements of cell density were made daily. It was often hampered
by the occurrence of condensation droplets on the cover of the 12-well
plates. To minimize light scattering during densitometry, the digitized
images of the cultures were recorded from beneath through a mirror. The
Bio-Rad video camera system Geldoc 1000 was used, with the "molecular
analyst" software for quantification. To record images of the plates,
the hood of the camera was rotated to a horizontal position. A mirror
was introduced through the open door of the hood, at 45° across the
horizontal light path, held by a special stand, which also maintained
the 12-well culture plate horizontally on a glass plate and held the Geldoc light source, upside down, 10 cm above the cultures. Full frame
images were captured in constant light conditions, i.e. white light being adjusted so that just a few saturated pixels appeared
in red in the empty optical field. In these conditions, a reproducible
relationship was found between the absorbance of wells in densitometric
scanning profiles of recorded images and the amount of cells in the
wells. Images of culture plates without cells were used as base-line
reference. They showed an uneven light background. This fact was
compensated by the subtraction of an adapted base line. Measurements
and base-line profiles were drawn using the rectangle tool of the image
analysis software. Base lines were drawn by joining the crest highlight
points in each image between the wells and then subtracted. This method was found more accurate and reproducible than the use of 24 grids or
the elliptical objects tools proposed by the software. When possible
and necessary, condensation was fully eliminated by replacement of the
cover in axenic conditions.
Fluorimetric Measurements of Iron Uptake in Culture
Wells--
Cells remained in their culture plates during fluorimetric
iron uptake measurements. Respiration of the cells was maintained by
aeration of the iron uptake medium in the return tubing from the
fluorimeter cuvette to the plate.
Prior to fluorimetric measurements, cells were drained by pressing a
plastic pipette on the bottom of the wells and by aspiration of the
culture medium through this leaky filtering contact. After three washes
and drainings with iron-free culture medium (MS-Fe), the well was ready
to be connected to a 7-ml circuit, filled with the iron uptake medium
IUM (IUM = iron-free culture medium (MS-Fe) supplemented with
Tris-maleate NaOH buffer (0.02 M), pH 5.8, 4·10
The drained cells received 1.5 ml of IUM and were connected to the 7-ml
fluorimeter circuit described above. Excitation was made at 488 nm
(slit: 1.5 nm) and emission recorded at 511 nm (slit: 2 nm). After
recording the initial fluorescence level in the absence of cells for
200 s, the culture well and cuvette circuit were connected, and
changes of fluorescence were recorded kinetically. The same
spectrofluorimeter was used for in vitro iron exchange studies.
Iron Binding Properties of CA
Thermodynamic Studies
pKa Determination of CA and
FeIIICAH--
The thermodynamic analysis of the
Fe(III)-calcein complexation equilibria required the determination of
the ligand deprotonation constants and an equilibrium study of the
ferric complexes as a function of the pH. The deprotonation constants
of the ligand have been determined by potentiometric (Fig.
1) and UV-visible spectrophotometric
titrations. Analysis of the titration curves yield the
pKa values: 3.09 (3), 3.67 (3), 5.29 (3), 9.57 (2) (from potentiometry)
and 12.3 (1) (from spectrophotometry). Numbers in
parentheses represent the S.D. in the last significant digit. The
pKa values of three out of four of the carboxylic
acids of the aminodiacetic moieties are too low to be determined from
potentiometric measurement. The potentiometric titration curve of a 1:1
iron(III):calcein solution (Fig. 1) exhibited a buffer region at
a = 3-4 (a is the number of moles of added
base per moles of ligand) over the pH range 5-7. Another buffer region
is found at a = 4-5 over the pH range 8-9, indicating
that the predominant species at pH < 8 is the monoprotonated
complex, FeIIICAH. Its pKa has been
determined to be 8.90 (5) from the analysis of the titration curve. The
precipitation of the complex below pH 5 precluded the use of the
titration data in refinement program. The complex stability constant
was thus determined from spectrophotometric competition. The UV-visible
spectrum of the FeIIICAH complex at pH 7.2 exhibits a band
at 500 nm with a shoulder at 600 nm (Fig.
2)
In Vitro Equilibria and Iron Exchange
Competitions--
Competitions were carried out in order to determine
the relative efficiency of calcein and EDTA, NTA, and citrate anions
with regard to iron chelation and were observed spectrophotometrically. The FeIIICAH solution was obtained by adding equimolar
amount of Fe(III) perchlorate to CA at low pH, followed by a 0.05 M Tris-Cl, pH 7.2, 0.05 M NaClO4
buffer. Equimolar amounts of FeIIICAH complex and tested
ligands were left in the dark for 4 days at 25 °C. The reverse
exchange was done using the tested iron complex in competition with
calcein. UV-visible light spectra were recorded when exchanges in both
directions had led to similar results. The spectra in Fig.
2A show (i) a complete transfer of iron from
FeIIICAH to EDTA; (ii) an exchange with citrate, too low
for quantification; (iii) a significant exchange with NTA allowing the
quantification of iron(III) distribution between the two competing
ligands and the calculation of thermodynamic constants. In addition,
similar experiments were made in the presence of 2, 5, 10, and 20 molar eq of NTA with respect to FeIIICAH. The metal displaced was
measured to be 33% in the presence of 20 NTA eq. The competition
equilibrium can be expressed by the following equations:
where FeNTA, CA, and NTA refer to all forms of FeIII
complexed with NTA, free CA, and free NTA, at equilibrium and pH 7.2, respectively. Here UV-visible and EPR Studies
The UV-visible spectrum FeIIICAH exhibits a phenol to
iron charge transfer band at 600 nm (
Differences in coordination were also shown by EPR spectra of
FeIIIcomplexes of CA, EDTA, and calcein blue at pH 7.2; EPR
spectra of calcein and calcein blue are similar, as shown in Fig. 3
(the differential spectrum is very close to the base line). They are both centered at g = 4.293, and they are dissymmetric. The
spectrum of FeIIIEDTA is symmetrical and centered at g = 4.272.
These results suggest a coordination with one aminodiacetic arm and the
phenolate group of the calcein, while the second aminediacetic arm
remains uncoordinated. Since its amine pKa has been determined to be 9.6 in the free ligand, it is still protonated at pH
7.2. Consequently, calcein is found to act as a tetradentate ligand
like calcein blue.
Then, the question is raised of the possible binding of a second iron
cation by the second aminodiacetic arm. During the spectrophotometric titration of CA by iron(III), a sharp and intense peak at 550 nm
increased up to the addition of 2 eq (Fig.
4).
EPR titration experiments were performed under the same experimental
conditions. Fig. 5 shows that the EPR
signal intensity reaches a maximum at one iron equivalent per ligand
and then decreases until it vanishes when 2 eq of iron have been added.
It can be assumed that two successive bindings of iron cation occur,
involving the two aminodiacetic arms. The diferric complex is EPR
silent revealing antiferromagnetic coupling of the two iron ions.
Further physiological experiments will be limited to conditions where
only one iron(III) cation is complexed.
Electrochemical Studies
CV was performed in a 5 mM aqueous solution of the 1:1
FeII and FeIIIcalcein complex in the presence
of 0.1 M Tris-Cl, pH 7.2, 0.1 M
NaClO4 as supporting buffered electrolyte. Upon scanning
toward the cathodic region of potentials, an irreversible
electrochemical system is seen on the CV curve (Fig.
6) at Epc = Quenching of Fluorescence after FeIII and
FeII Chelation
An equimolar aqueous solution (2.5·10 Kinetic Iron Exchange Studies (in Vitro)
Citrate is known to be very abundant in vivo (1-8
mM citrate compared with 2-100 µM iron in
root sap (24)) and serves as the main intercellular iron transporter in
plants. The affinity of citrate for iron(III) is characterized by a log
A stopped flow experiment of iron(III) complexation by calcein and
calcein blue has shown that the reaction is completed in 1 s (Fig.
8). CA fluorescence is therefore expected
to be a fast reporting tool of the iron(III) availability.
Kinetics of iron exchange between calcein and NTA or citrate have been
followed by measuring the variations of the fluorescence intensity of
the solution during incubation with a large excess of NTA or citrate
(10 or 6320 molar eq, respectively) in a buffer 40 mM Hepes
at pH 7.2, 150 mM NaCl, at 25 °C in the presence of 4·10 In Vivo Kinetics of Iron Nutrition by Plant Cell Suspensions
Because of the high sensitivity of fluorescent probes, real time
measurements of iron(III) uptake by plant cells have been carried out.
Radioactive techniques are generally used for milligram amounts of
plant material for nutritional studies. They require destructive assays
on aliquoted samples. Fluorescent assays instead are neither
destructive nor harmful and allow real time observations without
sampling. However, fluorimeter cuvettes are not adapted for plant cell
cultures and for discriminating signals coming from the cells from
those coming from the medium. A system in which plant cell growth can
be maintained was preferred, which allowed a separate measurement of
the fluorescence of the medium. Contrary to its acetoxymethyl ester
derivatives, unesterified CA is not able to cross the membranes and
remains extracellular. The apoplasm, the compartment mainly composed of
cell walls, is extracellular and is known to have an important iron
storage function in iron nutrition for plants (25-27). The device used
here has the advantage of measuring both intracellular and apoplastic
iron uptake. A [Fe]/[CA] = 1/2 mixture was used because it presents the strongest iron buffering power, and it can reflect both iron release and uptake. Culture plates were connected to the fluorimeter cuvette by a system allowing respiration with full oxygenation of the
medium. In some experiments, the ionophore A 23187 (Sigma) was added to
check if the cell membrane could be a limiting factor in the uptake
process. Fig. 10 shows that upon
contact with the cells, the fluorescence of the nutrient solution
raises regularly for more than 30 min, indicating iron release from
FeIIICAH. The rate of iron uptake by cells is thus directly
readable from the increase of fluorescence. The addition of the
ionophore accelerates the rate of the uptake. The sensitivity of the
method is high, and the rate of iron uptake is so high (up to 0.93 nmol·h
Calcein as a Fluorescent Probe for Ferric Iron
APPLICATION TO IRON NUTRITION IN PLANT CELLS*
,,,,
Laboratoire de Chimie Biomimétique,
UMR CNRS 5616, Université Joseph Fourier, Grenoble, France
B.P.53 38041, Grenoble cedex 9 France, the § Laboratoire
d'Electrochimie Organique et de Photochimie Rédox, UMR 5630, Université Joseph Fourier, Grenoble, France B.P.53 38041, Grenoble cedex 9, France, the ¶ Laboratoire de Chimie et Biochimie
des Centres Rédox Biologiques, CEA Grenoble/EP 1087 CNRS/Université Joseph Fourier, 17 rue des Martyrs, 38054 Grenoble cedex 9, France
ABSTRACT
TOP
ABSTRACT
INTRODUCTION
MATERIALS AND METHODS
RESULTS
DISCUSSION
REFERENCES
111 of FeIIICAH is 33.9. The free metal ion
concentration is pFeIII = 20.3. A
(FeIII)2 CA complex can be formed. A reversible
iron(III) exchange from FeIIICAH to citrate and
nitrilotriacetic acid is evidenced when these ligands are present in
large excess. The kinetics of iron(III) exchange by CA is compatible
with metabolic studies. The low reduction potential of
FeIIICAH shows that the ferric form is highly
stabilized. CA fluorescence is quenched by 85% after FeIII
chelation but by only 20% using FeII. Real time iron
nutrition by Arabidopsis thaliana cells has been measured
by fluorimetry, and the iron buffer FeIIICAH + CA was used
as source of iron. As a siderophore, FeIIICAH promotes cell
growth and regreening of iron-deficient cells more rapidly than
FeIIIEDTA. We conclude that CA is a good chemosensor for
iron(III) in cells and biological fluids, but not for Fe(II). We
discuss the interest of quantifying iron buffers in biochemical studies of iron, in vitro as well as in cells.
INTRODUCTION
TOP
ABSTRACT
INTRODUCTION
MATERIALS AND METHODS
RESULTS
DISCUSSION
REFERENCES
MATERIALS AND METHODS
TOP
ABSTRACT
INTRODUCTION
MATERIALS AND METHODS
RESULTS
DISCUSSION
REFERENCES
4 M calcein solution (0.05 M Tris-Cl, pH 7.2, 0.05 M NaClO4);
pH was controlled before and after each measurements. Iron
concentration was obtained spectrophotometrically by using a molar
extinction coefficient 
= 4160 M1·cm1 at 240 nm (20); CA
concentration was obtained by using the molar extinction coefficient
provided by Molecular Probes. The total volume was 500 µl, and its
variations due to iron additions from stock solutions were less than
2%. Equilibria after competitions were measured using the same equipment.
3 M
calcein, and increasing amounts of iron(III) perchlorate. Temperature was 100 K and field was scanned from 1000 to 5000 gauss.
) with early stationary phase
Arabidopsis thaliana (var. Columbia) cell suspensions. For
axeny, plates where sealed with "Magic Scotch®" tape. 24-well
plates were agitated at 220 rpm at 25 °C in a New Brunswik
"innova" 4230 refrigerated incubator shaker. Light was supplied
18 h a day by two growlux Sylvania 15-watt fluorescent tubes in
the "photosynthesis" accessory.
7 M CA, and 2·107
M ammonium iron(II) sulfate; the iron-CA complex was
synthesized previously at low pH and in a 100 times more concentrated
solution. It fully oxidizes upon dilution at pH 5.8. IUM is stored in
50-ml fractions at 
20 °C, protected from light, and used
immediately after thawing. An aspiration filter made with a 20-µl
pipette tip filled with glass wool is connected to a peristaltic pump. The medium is pumped out at a rate of 0.42 ml·min1 into
the cuvette of the fluorimeter. The rate of the return pumping was
twice as fast, therefore air was swallowed by the tubing, allowing both
the level control in the cuvette and an intimate aeration of the
medium. The filter was agitated to prevent cell stacking and to improve
aeration. The 400-µl cuvette of the fluorimeter (Jobin-Yvon Spec
Fluoromax) is connected to the circuit with two Teflon tubes.
RESULTS
TOP
ABSTRACT
INTRODUCTION
MATERIALS AND METHODS
RESULTS
DISCUSSION
REFERENCES
View larger version (11K):
[in a new window]
Fig. 1.
Potentiometric titration curves for 0.5 mM CA (a) and CA + FeIII 1:1 0.5 mM
(b). All solutions were at 25 °C and
I = 0.1 M (NaClO4). The complex tends to
precipitate in conditions indicated by the dotted line.
View larger version (19K):
[in a new window]
Fig. 2.
A, equilibrium competition of equimolar
amounts (1.5·10 4 M) of FeIIICAH, with EDTA,
or NTA or citrate, kept in the dark for 4 days at 25 °C pH 7.2 in
0.05 M Tris-Cl, 0.05 M NaClO4.
B, UV-visible spectra of FeIII complexes of
EDTA, calcein blue, and calcein; concentrations are
0.45·103 M, 0.1 M Tris-Cl
buffer, pH 7.2, 25 °C.
with K = ([FeNTA]tot[CA]tot[H+])/([FeCAH]tot[NTA]tot)·(
(Eq. 1)
NTA/(FeNTA·CA))
and K = (110FeNTA/111FeCAH)·(NTA/FeNTA·CA)),
's are the usual Ringbom coefficients (22)
calculated from the deprotonation constants of the ligands CA and NTA
(23) or from the complexation constants of FeIIINTA (23).
The average formation constant log 111 of
FeIIICA was determined to be 33.9(1). The typical
pFeIII value of calcein was calculated to be 20.3 (pH = 7.4, [Fe]tot = 1 µM,
[ligand]tot = 10 µM). It is thus clear that
FeIIICAH is almost a thousandfold weaker complex than
FeIIIEDTA(pFeIIIFeEDTA = 23.5).
= 900 M1·cm1), which is not present
in the FeIIIEDTA spectrum (Fig. 2). In order to gain
information about the mode of coordination of the calcein with Fe(III),
the spectrum was compared with that of the FeIIIcalcein
blue. Calcein blue is a fluorescent ligand having a single aminediacetic complexing arm (see Fig. 3)
instead of the two complexing arms of the diaminetetraacetic in EDTA
and calcein. The UV-visible spectrum of FeIIIcalcein blue
(Fig. 2B) exhibits a phenolate to iron charge transfer band
at 490 nm. FeIIICAH thus provides a coordination to iron
significantly different from EDTA unlike suggested previously (4).
View larger version (13K):
[in a new window]
Fig. 3.
Formulae and EPR spectra of
FeIIICAH (1),
FeIIIcalcein blue (2) and
FeIIIEDTA (4). Curve 3 is the
difference (1 2); 2' and
4' are expanded images of the major signal (1 mM
concentration, in 10% glycerol, 0.1 M Tris-Cl buffer, pH
7.2; T = 100 K).
View larger version (26K):
[in a new window]
Fig. 4.
UV-visible spectra of calcein progressively
saturated by up to 2.4 eq iron perchlorate. 10 4
M calcein, 0.1 M Tris-HCl buffer, pH 7.2, 25 °C. Arrows: iron/CA = 1 and
2.
View larger version (10K):
[in a new window]
Fig. 5.
Evolution of the intensity of EPR signals of
FeIIIcalcein along iron additions (1 mM concentration, in 10% glycerol, 0.1 M Tris-Cl buffer, pH 7.2; T = 100 K).
0.63
V versus SCE (0.38 V versus NHE). A shoulder is
also observed at the positive foot (0.54 V) (0.29 versus
NHE) of this main system. Irreversibility of the cathodic process
undergone by the FeIII complex suggests that the
corresponding FeII complex is not stable, i.e. a
strong change in the coordination sphere around the iron cation occurs
during electron transfer and/or the FeII cation is released
in solution as an aqueous hydroxo-complex under our experimental
conditions (pH 7.2). In addition, the low reduction potential of the
FeIII complex shows for the first time that the
FeIII redox state is highly stabilized upon complexation
with calcein. Compared with the free FeII/FeIII
system in acidic solution (Ef = 0.52 V
versus SCE) (0.77 V versus NHE),
FeIII is stabilized by ~1.15 V. If the overall stability
constant for FeIICAH (assuming a protonated complex as for
FeIII) log 111(FeII) is evaluated according
to the following formula, Ec° = Ef° + 0.059log[111
(FeIICAH)/111
(FeIIICAH)], its value is 14.4, that means that when the
reduction of the complex occurs, the FeII complex is
clearly partially dissociated.
View larger version (9K):
[in a new window]
Fig. 6.
Cyclic voltammetric curve of iron CA (5 mM) in Tris-Cl, pH 7.2, 0.1 M
NaClO4 aqueous solution; scan rate 0.1 V
s 1; working electrode: vitrous carbon disc electrode (5 mm diameter); E versus SCE.
7
M) of CA and Fe(III) is made. Similarly, FeII + CA mixture is prepared in a glove box. The fluorescence emission spectrum of CA is compared with that of Fe(III) and Fe(II)CA complexes (Fig. 7). Quenching was 85% with Fe(III)
and only 20% with Fe(II). At the same time, the UV-visible spectrum of
a CA + FeII solution at 104 M was
recorded and found to be very similar to that of CA alone (not
shown).
View larger version (14K):
[in a new window]
Fig. 7.
Quenching of fluorescence after addition to
CA of FeIII or FeII in equimolar amount.
Concentrations are 2.5·10 7 M, pH 7.2. a = CA, b = CA + FeII, and
c = CA + FeIII. b is prepared
using a glovebox in argon atmosphere.
110(LFeIII) value of 11.8. Thus, this value seems low as
compared with that of calcein which is measured here, log
111(LFeIII) = 33.9; one would expect calcein to take up
iron from citrate. Ideally, a fluorescent chemosensor should not
withdraw too much iron from iron transport molecules, which are the
precursor of most biological iron ligands. NTA (log
110(LFeIII) = 15.9) has an intermediate affinity, maybe
close to that of many cellular iron ligands. Experiments are necessary
to test if citrate has the potential to share iron with calcein (and
NTA) in an inversible equilibrium, which evolves quickly enough to
reflect metabolic competition for iron.
View larger version (11K):
[in a new window]
Fig. 8.
Kinetics of Fe(III) complexation by CA.
Stopped flow analysis of FeIII complexation by CA or
calcein blue ([FeIII] = 0.03 M,
[H+] = 0.03 M, [NaClO4] = 2 M, [calcein] = 2·10 5 M,
[calcein blue] = 2.5·104 M) temperature
is 25 °C, and absorbance is monitored at 522 nm for calcein and 530 nm for calcein blue; curves are normalized.
7 FeIIICAH. These conditions were
chosen because citrate is much more abundant in vivo as
compared with CA, when introduced as a probe. Fig.
9 shows that fluorescence intensity at
511 nm increases, indicating iron release from calcein. The kinetics of
iron release from FeIIICAH to NTA or to citrate is easily
measurable in the time scale of minutes. Calcein is shown able to
reflect both iron loading or release by citrate, in the range of
concentrations expected to be found in vivo. We conclude
that despite its high affinity, calcein at low concentrations seems
able to reflect changes in pFeIII induced by ligands like
citrate and NTA, which cover a wide range of affinity for iron,
presumably that of metabolic iron ligands.
View larger version (11K):
[in a new window]
Fig. 9.
Kinetics of the fluorimetric changes of
FeIIICAH 4·10 7 M
after the addition of either NTA 2.105
M or citrate 2.5·103
M at pH 7.2 in the presence of 40 mM Hepes, pH 7.2, and 150 mM
NaCl.
1/g of fresh cells) that it cannot be considered
as a steady state rate of iron nutrition. Therefore, CA is potentially
a good siderophore.
View larger version (9K):
[in a new window]
Fig. 10.
Fluorimetric measurement of the rate of iron
uptake by A. thaliana cells in suspension
culture. Growth medium is washed out and replaced by IUM medium
(2·10 7 M FeIII and
4·107 M CA). Iron uptake is measured by its
release from calcein. The ionophore A 23187 is added at 1200 s.
Growth and Greening of Plant Cells Using FeIIICAH as Iron Source
In order to test the physiological significance of the iron uptake observed using fluorescence monitoring, we have followed the growth of cell suspensions fed with FeIIICAH as the only source of iron. The inocula derived either from cell suspensions prestarved with iron (presenting limited symptoms of chlorosis) or from controls provided with sufficient iron.
Results are shown in Fig. 11. Using
iron-prestarved cells or normally fed cells, FeIIICAH was
found to be a very good iron source. In 7 days, growth and regreening
after chlorosis was much faster than when control FeIIIEDTA
was used. However, faster growth or greening were hardly noticeable
when, in other experiments (not shown), nonprestarved inocula were
used.
|
A special experiment was set up to test the physiological function of
the iron apparently taken up at high rate by cells during the initial
contact with FeIIICAH; we wanted to check that neither
precipitation nor other nonphysiological artifact interfered with iron
uptake. Inocula from iron-prestarved and iron sufficient cultures were
pretreated for 1000 s with 50 µM iron as
FeIIICAH or as FeIIIEDTA. A series was then
washed and cultivated without iron, another was maintained in the
presence of 50 µM iron as FeIIICAH or
FeIIIEDTA. Growth and chlorotic symptoms (Fig.
12 and Table
I) were followed along the cultures.
|
|
The growth of iron-sufficient inocula was not affected after a 8 days culture without iron (wells 13-15) but cells became chlorotic. Nevertheless, a pretreatment of 1000 s with FeIIICAH was sufficient to prevent a further appearance of chlorosis (wells 19-21), while after the same treatment using FeIIIEDTA, the cells became chlorotic (wells 22-24). The fast initial uptake had an efficient nutritional function.
Using prestarved cells, pretreatment of 1000 s with FeIIICAH limited chlorosis (wells 7-9 compared with 1-3), but did not restore normal growth. Pretreatment of 1000 s with FeIIIEDTA had no positive effect (wells 10-12). In controls, the continuous presence of 50 µM FeIIIEDTA restored growth and greening (wells 16-18). The same amount of FeIIICAH killed the inoculum (wells 4-6) during the lag time period, at high dilution of cells.
We conclude that FeIIICAH is a more readily available source of iron for plant cells than FeIIIEDTA, since it favors a faster growth and regreening (Fig. 11). The iron taken up from FeIIICAH at high initial rates, during the 1000 s treatments (or during fluorescence uptake measurements), can be physiologically mobilized later, and is therefore efficiently stored in cells or/and in the apoplasm. FeIIIEDTA does not allow such fast iron storage loading.
Nevertheless, toxic effects have been evidenced at
5·105 M FeIIICAH, only on iron
prestarved cells, in a dilute weakened inoculum, where the amount of
iron per cell is higher than in normal growth or during iron uptake
experiments. This toxicity may be correlative of an easy uptake of iron
and may be due to Fenton reactions after iron reduction. Some
photolability of FeIIICAH have been observed in
vitro (not shown). It may be due to photoreduction followed by
Fenton reactions as observed with FeIIIcitrate and
FeIIIEDTA (28).
| |
DISCUSSION |
|---|
|
TOP
ABSTRACT
INTRODUCTION
MATERIALS AND METHODS
RESULTS
DISCUSSION
REFERENCES
|
|---|
From a chemical point of view, calcein is an interesting iron ligand whose fluorescence is quenched upon binding to iron. As a consequence, it has been studied as a sensor for iron and suggested to be a useful tool for monitoring cytosolic iron and assessing the dynamics of intracellular iron in living cells; interesting observations were evidenced (4 to 10). Nevertheless, no firm characterization of its thermodynamic, dynamic, and structural properties had been carried out previously and this was the first aim of this work.
We show for the first time that calcein is a specific FeIII
ligand and binding of FeII is scarcely sensed by
fluorescence extinction. The affinity for FeIII is shown
from the large 111 value and the large
pFeIII value. This is not surprising since calcein is
providing in addition to carboxylates, a phenolate for binding iron
(29) as shown from the characteristic phenoxo-to-iron charge transfer
band ( = 900 M1·cm1 at 600 nm) in its visible spectrum. This band is also present in the spectrum
of the calcein blue-iron complex ( =1200
M1·cm1 at 490 nm). This
combination of carboxylates and phenolate is known to greatly stabilize
the FeIII form. This is reflected in the high stability
constant and the very low redox potential, much comparable with those
for siderophores. It is clear that calcein is significantly different
from EDTA as a ligand: different coordination (thus different visible
and EPR spectra), lower redox potential of the
FeIII/FeII couple. An unexpected property of
calcein is its ability to bind 2 FeIII atoms; the absence
of EPR signal and the very intense absorption band at 550 nm of the 2:1
iron:calcein complex are remarkable physicochemical characteristics.
From a biochemical point of view, another interesting property of calcein is its fast exchangeability as an iron ligand. This is reflected in the fast kinetics of iron complexation during incubation of FeIII and calcein and in the fast exchange with other iron(III) ligands, such as NTA or citrate the iron transporter in plants, over a wide range of affinities. Thus the calcein/iron-calcein system is an iron buffer sensitive to competition with other ligands. This makes it a kinetically competent, simple, sensitive, and nondestructive fluorescent probe for monitoring FeIII in cells.
Our finding that calcein is a poor fluorescent probe for
FeII is in marked contrast with conclusions from several
papers (4, 10). The redox conditions required to form a
FeIIcalcein are not physiologically relevant and thus any
binding, if any, of FeII will result in oxidation and
formation of the FeIIIcalcein complex, disrupting the
balance between FeII and FeIII. Our data thus
do not support the assumption that calcein fluorescence has the
potential to reflect the variations in the ferrous iron cytosolic pool.
Experiments using 55Fe and organic solvent extractions of
iron complexes were presented in Ref. 4 as corroborating the hypothesis
of a ferrous probe. They are not convincing since tissue grinding is
necessary, therefore cellular compartmentation is destroyed, oxygen is
introduced, redox potential and pH are averaged, and metabolic iron
fluxes are interrupted. No consensus appears in the literature on
estimations of cellular FeII concentrations: from
107 M or 108 M (10)
to 1017 M as estimated through the affinity
for nicotinamine, the FeII transporter in plants, which is
suspected to be the natural FeII sensor (30).
From a physiological point of view, we demonstrate two new applications of the FeIIICAH complex. We show (i) that it provides the first real time tool for monitoring iron(III) uptake by intact plant cells. The contribution of the apoplasm, which is physiologically important for iron storage (25-27), is not bypassed by the proposed method. Furthermore, (ii) a physiological study of iron-calcein as a nutritional siderophore demonstrates that CA is better than EDTA for restoring normal growth after iron starvation and chlorosis; its toxicity only appears in unusual high concentrations. The nutritional properties of FeIIICAH reflect its chemical affinity for iron: an intermediate between those of citrate and EDTA. Calcein is a fast source and the iron taken up for 1000 s from calcein is further functional for growth and prevention of chlorosis and therefore stored in a physiologically efficient compartment, presumably ferritins and the apoplasm. CA is therefore expected to bind iron in the range of affinity of the ferric cellular iron pool.
Iron buffering conditions are very often disregarded. In many in vitro experiments, iron is introduced either as a salt or as a more stable 1/1 iron complex. In these conditions, iron is not buffered and a small exchange of iron or ligand results in drastic changes in free iron concentrations. An appropriate excess of ligand should be introduced to buffer the iron donor system (as it is done for calcium standards (31)). pFe calculations are generally possible in vitro. They are directly predictive of a hierarchy in actual binding efficiencies. The quantification of metal buffering could help biologists to describe and compare their experimental conditions. Like pH is generally controlled in the experiments describing acido-basic reactions, buffered iron reagents (i.e. in known pFe limits), are desirable in experiments on iron metabolism.
In animal cells, the limit between iron deficiency and sufficiency is sensed naturally by iron regulatory proteins (IRP 1 and IRP 2), which are ligands for FeII. FeII loading of IRPs act as a molecular switch for iron uptake or storage. According to Epsztejn et al. (10), the definition of the LIP is the sum of free and CA bound iron. This definition invokes FeIII and is not based upon the iron buffering power of the cells. We propose a physiological definition of the LIP; the LIP is the amount of buffered iron that can be released by the cell before the IRPs loose their iron.
New fluorescent ligands whose affinity for iron(II) is close to that of
the IRPs must be characterized for the quantitative estimation of
cellular iron(II). Some fluorescent probes become available for
iron(III) (32, 33). A fluorescent form of desferrioxamine B has been
synthesized. It was not found able to measure directly iron nutrition
by plants, except when used in symbiosis by the microflora of the roots
(34). A ferrichrome analog has been used to measure iron uptake by
bacteria (35). A search for ferrous fluoroprobes is needed.
| |
Acknowlegments |
|---|
We thank Dr. A. Martre for technical assistance and Dr. A. Deronzier for his interest on this work.
| |
FOOTNOTES |
|---|
* The costs of publication of this article were defrayed in part by the payment of page charges. The article must therefore be hereby marked "advertisement" in accordance with 18 U.S.C. Section 1734 solely to indicate this fact.
To whom correspondence should be addressed. Tel.:
33-4-76-88-91-02; Fax: 33-4-76-88-91-24; E-mail:
laulhere{at}cbcrb.ceng.cea.fr.
2
pFe = log[Fe] = log 110 + log ([L]T FeL/L
[FeL]T), where 110 is the overall
stability constant for the Fe + L 
FeL equilibrium, [L]T = [L] +[FeL], [FeL]T = [Fe] + [FeL], and FeL and L are the Ringbom
coefficients (22). Chaberek and Martell (11) and others adopted this
logarithmic notation as convenient for low values and reminiscent of
the familiar pH scale. In the presence of free ligand, a metal
concentration is buffered. Its concentration (i.e. pFe) is
least affected by disturbances, which tend to add or to remove metal
ions from the medium. Raymond and co-workers (15) proposed conventional
condition to compare the efficiency of iron chelators.
| |
ABBREVIATIONS |
|---|
The abbreviations used are: LIP, labile iron pool; CA, calcein; SCE, standard calome electrode; CAH, monoprotonated calcein; NTA, nitrilotriacetic acid.
| |
REFERENCES |
|---|
|
TOP
ABSTRACT
INTRODUCTION
MATERIALS AND METHODS
RESULTS
DISCUSSION
REFERENCES
|
|---|
- Becker, R., Eberhard, F., and Manteuffel, R. (1995) Plant Physiol. 108, 269-275[Abstract]
- Crichton, R. (1991) in Inorganic Biochemistry of Iron Metabolism (Horwod, E., ed) pp. 131-172
- Fontecave, M., and Pierre, J. L. (1991) Biol. Met. 4, 133-135[CrossRef][Medline] [Order article via Infotrieve]
-
Breuer, W.,
Epsztejn, S.,
and Cabantchik, Z. I.
(1995)
J. Biol. Chem.
270,
24209-24215
[Abstract/Free Full Text] -
Breuer, W.,
Epsztejn, S.,
Millgram, P.,
and Cabantchik, Z. I.
(1995)
Am. J. Physiol.
268,
C1354-C1351
[Abstract/Free Full Text] - Cabantchik, Z. I., Glicktein, H., Milgram, and Breuer, W. (1996) Anal. Biochem. 233, 221-227[CrossRef][Medline] [Order article via Infotrieve]
- Breuer, W., Epsztejn, S., and Cabantchik, Z. I. (1996) FEBS Lett. 382, 304-308[CrossRef][Medline] [Order article via Infotrieve]
- Breuer, W., Greenberg, E., and Cabantchik, Z. I. (1997) FEBS Lett. 403, 213-219[CrossRef][Medline] [Order article via Infotrieve]
-
Glicktein, H.,
Breuer, W.,
Loyevsky, M.,
Konijn, A. M.,
Libman, J.,
Shanzer, A.,
and Cabantchik, Z. I.
(1996)
Blood
87,
4871-4878
[Abstract/Free Full Text] - Epsztejn, S., Kakhlon, O., Glicktein, H., Breuer, W., and Cabantchik, Z. I. (1997) Anal. Biochem. 248, 31-40[CrossRef][Medline] [Order article via Infotrieve]
- Chaberek, S., and Martell, A. E. (1959) Organic Sequestering Agents, pp. 174-215, John Wiley & Sons, New York
- Schubert, J. (1963) Iron Metabolism, an International Symposium, pp. 467-496, Springer-Verlag, Berlin
- Price, C. A. (1968) Annu. Rev. Plant Physiol. 19, 239-248
- Laulhère, J. P., and Alquier-Bouffard, A. (1969) Physiol. Veg. 7, 277-296
- Harris, W. R., Raymond, K. N., and Weitl, F. L. (1981) J. Am. Chem. Soc. 103, 2667-2675[CrossRef]
- Williams, R. J. P. (1982) FEBS Lett. 140, 3-10[CrossRef][Medline] [Order article via Infotrieve]
- Crumbliss, A. L. (1991) in Handbook of Microbial Iron Chelates (Winkelmann, G., ed), pp. 177-232, CRC Press, London
- Laulhère, J. P., Devillers, P., and Pierron, M. (1972) Physiol. Veg. 10, 589-598
- Gans, P., Sabatini, A., Vacca, A. (1985) J. Chem. Soc. Dalton Trans. 1195-1200
- Bastian, R., Weberling, R., and Palilla, F. (1956) Anal. Chem. 28, 459-462[CrossRef]
- Murashige, T., and Skoog, F. (1962) Physiol. Plant. 15, 473[CrossRef]
- Ringbom, A. (1963) in Complexation in Analytical Chemistry, Interscience, New York
- Martell, A., and Smith, R. M. (1974) Critical Stability Constants, Plenum Press, New York
-
Brown, J. C.,
and Chaney, R. L.
(1971)
Plant Physiol.
47,
836-840
[Abstract/Free Full Text] -
Zhang, F.-S.,
Römeld, V.,
and Marshner, H.
(1991)
Plant Physiol.
97,
1302-1305
[Abstract/Free Full Text] -
Longnecker, N.,
and Welch, R. M.
(1990)
Plant Physiol.
92,
17-22
[Abstract/Free Full Text] -
Pich, A.,
and Scholz, G.
(1991)
J. Exp. Bot.
42,
1517-1523
[Abstract/Free Full Text] - Caris, C., Baret, P., Béguin, C., Serratrice, G., Pierre, J. L., and Laulhère, J. P. (1995) Biochem J. 312, 879-885
- Saari, L. A., and Seitz, W. R. (1984) Anal. Chem. 56, 810-813
-
Stephan, U. W.
(1995)
J. Exp. Bot.
46,
531-537
[Abstract/Free Full Text] - Tsien, R., and Pozzan, T. (1989) Methods Enzymol. 172, 230-262[Medline] [Order article via Infotrieve]
- Fagès, F., Bodenant, B., and Weil, T. (1996) J. Org. Chem. 61, 3956-3961[CrossRef][Medline] [Order article via Infotrieve]
- Nudelman, R., Ardon, O., Hadar, Y., Chen, Y., Libman, J., and Shanzer, A. (1998) J. Med. Chem. 41, 1671-1678[CrossRef][Medline] [Order article via Infotrieve]
-
Bar-Ness, E.,
Hadar, Y.,
Chen, Y.,
Sanzer, A.,
and Libman, J.
(1992)
Plant Physiol.
99,
1329-1335
[Abstract/Free Full Text] - Weizman, H., Ardon, O., Mester, B., Libman, J., Dwir, O., Hadar, Y., Chen, Y., and Shanzer, A. (1996) J. Am. Chem. Soc. 118, 12368-12375[CrossRef]
Copyright © 1999 by The American Society for Biochemistry and Molecular Biology, Inc.
CiteULike 
Complore 
Connotea 
Del.icio.us 
Digg 
Reddit 
Technorati What's this?
This article has been cited by other articles:
|
|
 |
|
  Y.-S. Sohn, W. Breuer, A. Munnich, and Z. I. Cabantchik Redistribution of accumulated cell iron: a modality of chelation with therapeutic implications Blood, February 1, 2008; 111(3): 1690 - 1699. [Abstract] [Full Text] [PDF]
|
|
|
|
 |
|
 K. Takano, Y. Kitao, Y. Tabata, H. Miura, K. Sato, K. Takuma, K. Yamada, S. Hibino, T. Choshi, M. Iinuma, et al. A dibenzoylmethane derivative protects dopaminergic neurons against both oxidative stress and endoplasmic reticulum stress Am J Physiol Cell Physiol, December 1, 2007; 293(6): C1884 - C1894. [Abstract] [Full Text] [PDF]
|
|
|
|
|
|
M. Shvartsman, R. Kikkeri, A. Shanzer, and Z. I. Cabantchik Non-transferrin-bound iron reaches mitochondria by a chelator-inaccessible mechanism: biological and clinical implications Am J Physiol Cell Physiol, October 1, 2007; 293(4): C1383 - C1394. [Abstract] [Full Text] [PDF]
|
|
|
|
 |
|
 G. Liu, D. L. Greenshields, R. Sammynaiken, R. N. Hirji, G. Selvaraj, and Y. Wei Targeted alterations in iron homeostasis underlie plant defense responses J. Cell Sci., February 15, 2007; 120(4): 596 - 605. [Abstract] [Full Text] [PDF]
|
|
|
|
 |
|
 U. RAUEN, F. PETRAT, T. LI, and H. DE GROOT Hypothermia injury/cold-induced apoptosis--evidence of an increase in chelatable iron causing oxidative injury in spite of low O2-/H2O2 formation FASEB J, October 1, 2000; 14(13): 1953 - 1964. [Abstract] [Full Text]
|
|
|
|
 |
|
 A. M. Romeo, L. Christen, E. G. Niles, and D. J. Kosman Intracellular Chelation of Iron by Bipyridyl Inhibits DNA Virus Replication. RIBONUCLEOTIDE REDUCTASE MATURATION AS A PROBE OF INTRACELLULAR IRON POOLS J. Biol. Chem., June 22, 2001; 276(26): 24301 - 24308. [Abstract] [Full Text] [PDF]
|
|
| ||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
| HOME | HELP | FEEDBACK | SUBSCRIPTIONS | ARCHIVE | SEARCH | TABLE OF CONTENTS |
| All ASBMB Journals | Molecular and Cellular Proteomics |
| Journal of Lipid Research | ASBMB Today |