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Studies on the transport, metabolism and chemistry of iron-sugar chelates

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Studies on the transport, metabolism and chemistry of iron-sugar chelates

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Content STUDIES ON THE TRANSPORT, METABOLISM AND
CHEMISTRY OF IRON-SUGAR CHELATES
By
Bibudhendra Sarkar
A Dissertation Presented to the
FACULTY OF THE GRADUATE SCHOOL
UNIVERSITY OF SOUTHERN CALIFORNIA
In Partial Fulfillment of the
Requirements for the Degree
DOCTOR OF PHILOSOPHY
(Bi oohemistry)
August 19&+
U N IV E R S IT Y O F S O U T H E R N C A L IF O R N IA
T H E G R A D U A TE S C H O O L
U N IV E R S IT Y PA RK
L O S A N G E L E S , C A L IF O R N IA 0 0 0 0 7
This dissertation, written by
....... Bjlbudhejidra.-Satitar.... ....
under the direction of / t l . 8 ....Dissertation C o m
mittee, and approved by all its members, has
been presented to and accepted by the Graduate
School, in partial fulfillment of requirements
for the degree of
D O C T O R O F P H I L O S O P H Y
D ean
D a te Angus! 196U
DISSERTATIO N C O M M ITTEE
, —V
> C hairm an
ACKNOWLEDGMENT
It is with gratitude that I acknowledge the
guidance, advice and support of a brilliant teacher and
able research leader, Dr. Paul Saltman.
Special thanks are also due to the members of my
Doctoral Committees Dr. John Mehl, Dr. Arvan Fluharty and
Dr. Arthur Adamson. Their unique abilities and Interests
in physical and biological chemistry provided many ideas
for overcoming theoretical and experimental difficulties.
I offer my since rest thanks to Dr. Sidney Benson
for many helpful discussions and criticisms during my
graduate work.
Without the help of my friends and co-workers
Clyde Stitt and Harold Helbook, this research would have
been Immeasurably longer and definitely much less enjoy
able .
I would also like to express my appreciation to
the Faculty of the Department of Biochemistry and to my
fellow students for their many kindnesses during my
graduate career.
A special offer of thanks goes to Eleanor James
whose very kind help and advice throughout my years of
graduate study have made the most difficult aspects of
this endeavor much more pleasant.
ii
TABLE OP CONTENTS
LIST OP TABLES
Pag®
iv
LIST OP ILLUSTRATIONS
LIST OP ABBREVIATIONS vii
Chapter
I. HISTORICAL INTRODUCTION 1
Control and Regulation of Iron Absorption
Iron Transport and Metabolic Energy
The Role of Transferrin
The Role of Ferritin
Chelates in General-Chemistry of Metal
Chelates
Chelation and Iron Absorption
Chelates in Biological Systems
II. OUTLINE OP PROBLEM AND METHOD OP ATTACK .... 22
III. EXPERIMENTAL AND RESULTS...................... 2*+
Regulation of Iron Absorption and
Utilization
Chemical Characterization of Iron-Sugar
Complexes
The Stability Constant of Iron/t h )"Fructose
at Aoid pH
Partied Periodation of Iron-Sugar Complexes
IV. DISCUSSION................................... 85
V. SUMMARY..................................... 96
LITERATURE CITED.................................... 100
iii
LIST OP TABLES
Table Page
1. Percentage of the total iron transported in
tvo hours per gram of wet tissue for the
various chelates tested ................... 33
2. The distribution of iron in various organs
after in vivo incubation................... 31 * -
3. The distribution of iron in various organs
after in vivo incubation using NTA chelate
with and without the prior addition of
apotransferrin to circulating blood........ 35
*f. Distribution of radioactivity in the contents
of the intestinal segment................. 39
5. Recovery of radioactivity after periodation
of labeled glucose and ferric glucose .... 81
6. Distribution of radioactivity in 100/ul of
final periodate reaction product of labeled
glucose and ferric glucose............... 83
iv
LIST OP ILLUSTRATIONS
Figure Page
1. "Mucosal block" hypothesis for regulation
of iron uptake.............................. 2
2. Apparatus to study in vitro transport
of iron...................................... 26
3. Percentage of the total iron transported per
gram of tissue for the various chelates
tested...................................... 30
M-. Rate of iron transport in vitro as a function
of concentration of iron citrate............ 31
5. Kinetics of relative incorporation ofQradio-
activity into liver slices from Fe^-trans
ferrin ...................................... **3
6. Kinetics of relative incorporation of radio
activity from Fe^-chelate into liver slices . U-5
7. Difference spectrum of ferric fructose com
pared to ferric i o n ........................
8* Titration curves of Fe^+ and Fe^+ separately
and together in the presence of excess
fructose....................... *+9
9. Redox potential as a function of pH with and
without fructose........................... 51
10. Percentage equilibrium reached as a function
of time in the dialysis of Pe-chelates • • . • 5*
11* Infra-red spectra of fructose and ferric
fructose............. 58
12* The absorbancy at 370 nuu of solutions contain
ing 2.5 x 10“* M and 5 x 10“* (Fe)t at
increasing concentrations of fructose at
pH 2.5...................................... 68
v
Figure Page
13* The absorbancy at 370 mp of a 1.0 M fructose
solution in the presence of increasing
concentrations of Fe3+ at pH 2*5............ 69
I1 ** The graphic presentation of the data from
Figure 12................................... 70
15. Hate of change of optical density due to
change in concentration of periodate on
oxidation of glucose measured at 300 mju
with an Initial pH of 9 . 0 .................. 77
16. Apparatus for partial periodation.............. 79
17. Proposed scheme for the mechanism of iron
absorption from the Intestine by chelation . . 88
18. Proposed structures of ferric fructose
chelate..................................... 93
vi
LIST OP ABBREVIATIONS
1.
BAETA
— —
Bis 2-aminoethylethertetraacetic acid
2. BAL —
2, 3“dimercaptopropanol
3.
Dime don
—
Dime thyldihy drore so rc inol
DTPA
—
Diethylenetriaminepentaacetic acid
5.
EDTA
—
Ethylenediaminetetraacetic acid
6. EHPGr
—
N, N'-ethylenebis -2-(0-hydroxyphenyl)-
glyc ine
7.
Hyamin
—
p-(Diisobutyl-cresoxyethoxyethyl) dimethyl
benzylammonium hydroxide
8. Nembutal
—
Sodium pentobarbital
9.
NTA
—
Nitrilotriacetic acid
10. POPOP
—
1, l f-Bis-2-(5-phenyloxazolyl)-benzene
11.
PPO
—
2, 5”Diphenyloxazole
vii
CHAPTER I
HISTORICAL IRTRODUCTIOIT
Control and Regulation of
Iron Absorption
The most widely accepted mechanism for the regula
tion of iron transport from the lumen of the gut into the
blood is the Mmucosal block" as originally proposed by Hahn
and associates (1) and elaborated upon by Cranick (2). The
salient features of the scheme are presented in Pigure 1.
This hypothesis assigned to the intestinal mucosa the regu
latory ability to permit iron absorption in time of need
and iron rejection when the body’s needs are fulfilled.
The essence of the mucosal block theoiy is that ferrous
ions entering the mucosal cell stimulate the production of
an acceptor protein apoferritin. After entry into the
cell, the divalent ions are oxidized to the ferric state
by the redox potential at the luminal border, and the
trivalent iron combines with apoferritin. At the vascular
border of the cell the ferritin is reduced by a different
redox environment with consequent release of ferrous ions
which emerge from the cell into the blood stream and after
reoxldation are bound to transferrin, the iron transporting
1
2
LUMEN EPITHELIAL CELL PLASMA
Fe+++
on 1
Fe++ —
Apoferritin Apoferritin
|>---»Ferritin..—1 (H)
Pe+++
(0) t 1
—► Fe++ Fe++ —
Transferrin
) ---->
Fe+++
(o) t
—► Fe++
Figure 1.— "Mucosal block" hypothesis for
regulation of iron uptake.
globulin of plasma. Thus, when the needs of the body are
met, further absorption of iron is curtailed by the accumu
lation of mucosal ferritin. An additional pathway is
included for condition of exceptional iron need when the
redox levels of the cell allow ferrous ions to traverse
the cell directly.
The experimental evidence that led to Granick's
mucosal block theory is meager. One of the first observa
tions was made on iron-depleted dogs fed 100 mg ferric iron
followed by an oral test dose of radioactive iron several
hours later (3). The absorption of the test dose of radio
active iron was decreased or "blocked" when preceded by a
large dose of unlabeled iron. Evidence suggesting that
ferritin is Involved as the acceptor substance in the
mucosal wall was obtained in the studies on guinea pigs (*f)
and later on horses (5). Following the oral administration
of large amounts of iron, ferritin crystals could be seen
to accumulate in the mucosa to a maximum level after six
hours. These disappeared slowly during the next twenty-
four hours, paralleling the duration of the block to iron
absorption in dogs. These slender threads of evidence con
stitute the basis for a theory that has been accepted
widely for almost two decades.
In evaluating "mucosal block" theory, the following
points should be considered: Evidence has never been
I f
brought forward to support the theory that iron in the
ionic form penetrates the mucosal cell surface. Different
redox levels at the luminal and vascular borders of the
oell have never been measured, nor has any change of cell
redox levels in anemia been established. Mucosal ferritin
accumulation, except in the form of an iron storage com
pound, has not been found in the dog, man, rat or several
other species (6). Moreover, recent experiments cast doubt
on the mucosal block theory. With the administration of
increasing amounts of iron, the total amount absorbed is
increased even though the percentage absorption is dimin
ished (7, 8).
In recent years there has been an increasing number
of experiments concerning iron metabolism which cannot be
reconciled with the mucosal block hypothesis. The compre
hensive review of Moore (9) succinctly outlines many com
pelling reasons for the abandonment of the "mucosal block"
hypothesis as the fundamental mechanism for the regulation
of intestinal iron transport. Saltman and co-workers (10)
have shown by kinetic studies that both Fe++ and Fe+++
can be absorbed by the rabbit, if these ions are presented
as soluble uncharged complexes of low molecular weight. As
long as the iron remains available in a suitable chelate
form, the metal will continue to be absorbed up to the
limits of transferrin capacity.
Moat studies of iron absorption have utilized
ferrous iron (11) since it had been concluded that iron fed
as ferrous iron is better utilized than iron fed as ferric
iron (12)• It has been demonstrated that ascorbic acid
(13) and sulfhydryl compounds (I1 *) enhance iron absorption,
assertedly due to the fact that they are able to keep the
iron in the ferrous form. Groan, et al. (15) found that
several organic and amino acids including glutamic, aspar
tic, citric and tartaric as well as ascorbic increased iron
uptake. They ascribed this enhancement to increased acid
ity in the gut which would prevent the precipitation of
ferric hydroxide. Ferric hydroxide precipitates unless the
pH is substantially below 5, and at pH 8, the maximum eon-
—18
centration of ferric ion possible is 10 M.
Divalent iron is far more soluble than trivalent
and this fact alone could explain its increased rate of
absorption. It is interesting to note, however, that in
cases of iron deficiency anemia, practically any inorganic
iron preparation can be used successfully (16). These
observations, in general, were obtained from clinical
trials where an increase in the circulating hemoglobin was
taken as evidence of satisfactory utilization (17, 18).
However, there is no evidence that ferrous or ferric ion Is
absorbed as such. Critical experimental data have not been
presented to differentiate between the relative efficiency
6
of the two soluble ionic forms of iron and of their many
possible organic and inorganic combinations.
The work of Herndon, et a.1. (19) has shown that
sorbitol enhances ferric iron absorption. Sorbitol is not
readily metabolized and has no reducing effect on trlvalent
iron. These authors postulate that carbohydrates and
polyols either stimulate the absorption mechanism of the
intestinal wall or protect various substances in the intes
tinal lumen.
In recent years, Interest has been focused on the
regulatory mechanism of the mucosa at the cellular level.
Conrad and Crosby (20) have shown that mucosa of the normal,
iron-replete animal rejects most of the dietaiy iron, but
accepts some which is needed and some which is not. The
unneeded radioactive iron remains in the epithelial cells,
thereby tagging them. The tagged cells are gradually
pushed to the tip of the villus— and then off--by the
advance of newly formed, untagged cells.
Many experiments suggest that the intestinal
mucosa's absorptive capacity for iron is controlled by
erythropoietic activity and iron stores (21, 22). Recently
Crosby, et al. (23) have shown that a diet which is
deficient in iron causes a rapid, marked increase in iron
absorption in rats. The increased absorption occurs in the
absence of a significant change in iron stores as judged by
7
the effect of an equivalent change In stores produced toy
phlebotomy, and in the absence of increased erythropoietic
activity as judged by the rate of removal of Fe^ from the
plasma. Mendel (2*0 has studied the possible role of bile
in the control of iron absorption.
Wheby and Crosby (25) showed that iron absorption
follow8 two phases, an initial period of rapid absorption
lasting up to two hours during which 60 to 80 per cent of
total absorption into the blood took place. The remainder
took place at a slow rate over the subsequent twelve to
twenty hours. Thus regulation of iron absorption appears
to involve two steps, mucosal cell uptake and subsequent
transfer to the blood. Recently Brown and co-workers (6)
have shown that intestinal iron uptake in rats involves two
processes: (1) an early rapid uptake of iron bound to
serine and possibly other amino acids, and (2) a subsequent
slow release form of iron bound to a non-ferritin protein.
Brown, et al. (6) studied the uptake of radioactive
iron by the mucosa of the duodenum of normal rats, and its
subcellular distribution. The soluble fraction of the
homogenized intestinal mucosa was fractionated into protein
and non-protein components, and the iron-binding materials
were isolated. The purest batch of mucosal iron-binding
protein was heterogenous but had maximum molecular weight
far below that of apoferritin or of ferritin complexes.
Although the non-protein fraction of the mucoea included
sugars, nucleosides, nucleotides, and Kreb's cycle inter
mediates, they showed no detectable radioactivity migrating
with any of these components. Most of the radioactivity
moved with two amino acids, serine and glycine. Occasion
ally a third amino acid, with the mobility characteristics
of valine, was found, although this was variable. Although
ferric-glycine complexes have been well characterized
(26, 27), and ferroglycine sulphate complexes have been
employed (28 , 29) in an unsuccessful attempt to enhance
iron absorption, the finding of iron associated with this
amino acid and/or Berine in the intestinal mucosa was
unexpected.
Several workers have reported enhanced absorption
of iron when cysteine, glutamic acid, or aspartic acid
(30, 15) has been added to the intestinal loops of rats or
cats. It was suggested that formation of complexes between
iron and these amino acids was responsible for the
increased iron absorption, although there was no demonstra
tion of the complex within the intestinal mucosa.
Iron Transport And Metabolic Energy
While considerable advances in the understanding of
the regulation and absorption of iron through the mucosal
wall have been made, there is no agreement as to the actual
process of transport. Saltman and co-workers (31“33)
working with rat liver slices, have shown uptake of iron by
the slices against apparent concentration gradients. The
iron offered to the cells as ferric ammonium citrate, moved
at a rate controlled by diffusion. In vitro, the iron
initially diffusing into liver cells concentrates in the
nucleus (3*0. These experiments were performed using iso
lated liver cells incubated in medium approximately 30 /*g
of iron (as ferric ammonium citrate) per ml.
Jacobi and his associates (30) claimed to have
demonstrated an active transport mechanism for iron across
the jejunum of the cat and the guinea pig. They failed,
however, to take into account the binding of plasma iron
to transferrin, when concentration gradients from gut lumen
to plasma were calculated in these preparations in which
the blood supply of the intestine was left intact. Further
more, it is apparent from their data that more iron was
transferred from the mucosal surface to serosal surface
when their in vitro preparations were incubated under
nitrogen than when incubated under oxygen.
Recently Dowdle, Schachter and Schenker (8) have
claimed that the movement of Fe^ from mucosal to serosal,
surface against the concentration gradient is dependent on
the process of active transport. Everted gut sacs prepared
from segments of the proximal small intestine of rats
5q
transported Fe 7 from the mucosal to the serosal surfaces
10
against concentration gradients in vitro. They found
active transport process to be maximal in the region of the
small intestine immediately distal to the pylorus and
diminished In the more distal segments. Addition of
ascorbic acid to the incubation medium markedly Increased
the active transport of Fe^ in vitro. However, working
with the same system of everted gut loop Brown and Justus
(35) showed that the movement of iron across the mucosa did
not depend on active transport. No explanation for these
conflicting experiments has been offered.
The Role of Transferrin
The plasma is the chief medium for iron transport
(36, 37). The theory of the mucosal block (l*f) suggests
that iron moves from the mucosal cell as ferrous ion into
the blood stream. The high oxygen tension in this environ
ment quickly oxidizes the iron to the ferric form. As the
trivalent ion, it is tightly bound to /3-^-globulin which
has been termed "transferrin” or "siderophillin” (38).
Wallenlus (39) bas shown that all measurable serum iron is
bound to this globulin so long as the iron binding capacity
of the protein is not exceeded. Normal human serum con
tains approximately lOO^ug of iron per 100 ml, with a
saturation value of approximately 300^ug per 100 ml.
Griblet, Hickman and Smithies identified several
genetic variants of transferrin by differences in
11
electrophoretic mobility (*f0) • Kinetic homogeneity of
iron-transferrin is assumed but physiological demonstration
of this has been lacking. Recently Dem, Alberto and Glynn
(»+l) using two isotopes of iron studied the characteristics
of the disappearance of radio iron from plasma. It was
found that transferrin iron is not kinetically homogeneous.
Several workers have suggested that iron absorption
is independent of the relative saturation of the iron bind
ing protein. Pollack, et al. (*+2) showed that an infusion
of Fe-salts into the artery supplying an isolated duodenal
loop of a dog in situ failed to stop the absorption of Fe
from the lumen of the gut. Hallberg and Solvell (*+3), pro
posed the dependence of absorption on iron binding capacity
of transferrin. But Pollack, et al. (^2) indicated that
the absorption of iron from the isolated loop continued
despite saturation or near saturation of transferrin's iron
binding capacity. The negligible influence of saturating
transferrin on iron absorption is perhaps, not unexpected.
Hemochromatotics continue to absorb iron despite near
saturation of their total iron binding capacity (Mf). More
over, Heilmeyer, et al. (*f5) studied a case of atrans-
ferrinemia which absorbed excessive iron.
Recently the association constant has been shown to
be about 10^ (**6). This iron-protein complex is very
sensitive to pH. Shade, et al. (^7) were able to follow
12
the dissociation of the complex by measuring the charac
teristic absorption at *f60 myu. They found the highest
stability in the alkaline region above pH 7.5. As the
solution is made more acid, dissociation takes place until
at pH i*.8 only half the iron is bound. At pH *+.0 there was
no spetropho tome trie evidence of any complex present.
Transferrin concentration is regulated by a multi
plicity of factors. It has been shown that the protein
Increases during enhanced requirement for iron as in
pregnancy and iron deficiency anemia when the total serum
iron is low. Conversely, the concentration decreases dur
ing hemochromatosis and transfusion siderosis, when the
total serum iron is high. It is also regulated in part by
factors controlling other plasma proteins, such as albumin,
during infections and allied conditions (^8) •
Jandle, et al. C^9) have shown that iron is effec
tively transferred to reticulocytes and incorporated into
heme at lower iron saturation levels of transferrin than
could be demonstrated using a similar system where rat
liver slices were competing for the transferrin iron.
These authors have shown that specific receptor sites exist
on or in the reticulocytes which bind the iron with great
avidity prior to incorporation into hemoglobin.
13
The Role of Ferritin
Most of the iron diffusing into the liver cell is
bound as cytoplasmic ferritin, although hemosiderin is also
present, especially when body stores are high. The mechan
ism by which iron is bound to ferritin has been investi
gated by several workers. Michael is (50) has stated that
while iron may be removed from ferritin without difficulty,
the reconstitution of ferritin from apoferritin could not
be accomplished. Neither iron salts in the ferric or
ferrous form, nor colloidal ferric hydroxide were effective
in reconstituting ferritin. However, Bielig and Bayer (51)
have reported that they were able to obtain crystalline
ferritin by oxidizing ferrous ammonium sulfate in the
presence of the apoferritin. Loewus and Fineberg (52) also
stated that they were able to make ferritin from apoferri
tin by incubation of the apoferritin with ferric ammonium
citrate and reducing agent or with ferrous ammonium sulfate
in air. In no case was the characterization of the recon
stituted ferritin complete.
Mazur and Carleton (53) suggested that ferritin
plays an active role as an Intermediary between iron origi
nating from plasma and heme synthesized by marrow and
reticulocytes. They incubated rat and human reticulocytes
59
in plasma with serum bound Fe . The radio-iron was incor
porated into ferritin and heme. Similar results were
The Role of Ferritin
Most of the iron diffusing into the liver cell is
hound as cytoplasmic ferritin, although hemosiderin is also
present, especially when body stores are high. The mechan-
ism by which iron is bound to ferritin has been investi
gated by several workers. Michaelis (50) has stated that
while iron may be removed from ferritin without difficulty,
the reconstitution of ferritin from apoferritin could not
be accomplished. Neither iron salts in the ferric or
ferrous form, nor colloidal ferric hydroxide were effective
in reconstituting ferritin. However, Bielig and Bayer (51)
have reported that they were able to obtain crystalline
ferritin by oxidizing ferrous ammonium sulfate in the
presence of the apoferritin. Loewus and Fineberg (52) also
stated that they were able to make ferritin from apoferri
tin by Incubation of the apoferritin with ferric ammonium
citrate and reducing agent or with ferrous ammonium sulfate
in air. In no case was the characterization of the recon
stituted ferritin complete.
Mazur and Carleton (53) suggested that ferritin
plays an active role as an intermediary between iron origi
nating from plasma and heme synthesized by marrow and
reticulocytes. They incubated rat and human reticulocytes
59
in plasma with serum bound Fe . The radio-iron was incor
porated into ferritin and heme. Similar results were
lb
obtained in vivo using rat marrow, and incorporation of
both ferritin and heme was enhanced by phenyl hydrazine
and depressed by bacterial endotoxin. After continued
incubation, ferritin-Fe^ declined and heme-Fe^ rose
slightly.
It has been postulated (5*0 that the ferric iron
in ferritin is reduced to the ferrous form before leaving
the surface of the protein. Mazur (55) has outlined a
scheme which proposes that ferritin normally has a pre
ponderance of disulfide groups and has ferric iron at the
surface of the molecule. In the hypoxic liver, ferritin is
converted to the active form in which free sulfhydryl groins
chelate the ferrous iron. Green and Mazur (56) have impli
cated uric acid as the compound responsible for the reduc
ing action.
Such conclusions are based on the rapid release of
iron by ferritin in the presence of reducing agents and by
the apparent inability of transferrin to bind ferric ion
from solution. It should be recognized, however, that it
is not possible to have appreciable concentrations of
ferric ion at hydrogen ion concentrations which will not
denature most proteins. Experiments involving ferric ion
at neutral pH require the use of a chelating agent to keep
the iron in a soluble form. It is possible that the sta
bility constant and dissociation rates of these ferric
15
chelates have affected the observed results and the con
clusions that have been drawn from them.
Chelates in General— Chemistry of Metal Chelates
Cyolic compounds in which a metal is co-ordinated
with two or more donor groups of a single molecule have
exceptional stability and in many cases have remarkable
and valuable properties. This type of ring formation was
named chelation by Morgan and Drew (57) and the term has
been used to classify all types of ring systems with
metals, without regard to the nature of the chemical bond
involved.
Chelates have long been recognized to play key
roles in the metabolism of many trace metals. The high
stability constant of transferrin iron is considered
indicative of a chelate ring participating in its struc
ture (58), The structures of the important metallopor-
phyrins have been reported (59)* Ferric hydroxide micelle
of ferritin are believed to be composed of iron atoms bound
to each other with oxygen bridges (60). Biological com
pounds such as amino acids and proteins, pterins, ribo
flavin, purines, nucleic acids and ascorbic acid are known
to form chelate rings (6l, 62).
The atoms which are capable of donating electrons
to metal ions— principally oxygen, nitrogen and sulfur— are
involved in the formation of chelates. Small changes in
16
pH can exert marked changes in the avidity with which the
metal is bound) depending upon the chemical nature of the
functional group of the ligand. Minimal changes in pH as
well as the concentrations of other metal ions can affect
transfer of an ion from one compound to another (6l). Many
transition meted salts are insoluble at physiological pH
conditions. However, there are numerous ligands distrib
uted throughout biological systems which can keep them in
soluble complex form.
The dependence of the stability constant of a
chelate on hydrogen ion concentration has been previously
noted. An additional requirement is a reasonable concen
tration of the ligand with respect to the meted ion and a
concomitant dearth of competing co-ordinating groups. If
satisfactoiy conditions are met and the pH is slowly
adjusted from an acidic to a basic condition, a chelate
will form. Indeed, this is one criterion for the demon
stration of the existence of a chelate (63). If the metal
ion forms an insoluble hydroxide, the rapid addition of
base, even in the presence of excess ligand, can cause a
precipitate. The presence of a sufficient amount of ligand
may then cause slow solubilization again.
In addition to the ability of organic ligands to
solubilize metal ions under basic conditions, the effective
charge in the vicinity of the ion is altered by the
17
displacement of one or more hydrogen ions from the organic
molecule. For example, ferric citrate (64, 65) can have a
net charge of +3* +1» 0, -1 or -2, depending on the pH.
The stability constants of these various forms are vastly
25
different and vary from less than 1 to 10 . In recent
years valuable information has been obtained regarding the
structures of the metal chelates by X-ray technique (66),
rotatory dispersion (67, 68), magnetic measurements (69)*
electron spin resonance technique (46) as well as measure
ments of visible and ultraviolet spectra (70, 71).
Chelation and Iron Absorption
The effect of chelation of iron on its absorption
from the alimentary tract has not been altogether clear.
Some reports (72) suggest probable enhancement while others
indicate inhibition (73) of iron assimilation when a
chelating agent is used. It has been suggested that iron
may traverse the intestinal wall while chelated by ascorbic
acid or certain amino acids and that such iron chelates
pass unchanged through the gut wall from the lumen to the
blood (74). Brown, et al. (6) suggested that Fe^-amino
acid complexes are formed in its intestinal lumen prior to
their entry into the mucosal cell.
Saltman and co-workers (10) have shown that both
Fe++ and Fe++,f are rapidly absorbed when presented to the
mucosal wall suitably chelated with fructose. Hopping (75)
18
Indicated that some chelates do traverse the mucosa intaot
while others may enter the cell but then dissociate.
Nlssim (76) studied urinaxy excretion and diffusibility of
different iron preparations. It was demonstrated that the
urinary iron excretion in the first twenty-four hours after
the intravenous injection of 22.5 mg/Kg of iron as a ferric
hydroxide-ferrous ascorbate complex amounted to 27A per
cent of the dosage. Foye, et al. (77) suggested that che
lation is a necessary requirement for the distribution and
utilisation of iron, since transport of ionic iron, per se,
through cellular membrane, is questionable. Recently Rubin
and Princiotto (78) reviewed the roll of chelation as a
basic biological mechanism. They showed that the reaction
of metal ions and organic ligands provided a buffer system
for the concentration of metal ion in the medium. Modifi
cation of the ligands by intermediary metabolism or meta
bolic alterations may selectively vary the metal ion con
centration.
Chelation in Biological Systems
Nisland8 (79) has shown that some fungi and bacteria
secrete several strong iron chelates called •ferrichromes'
which are important in their iron metabolism. Organic
acids are believed to play an important role in the trans
port of trace metals in plants. Recently Tiffin and
Brown (80) have shown that malonlc acid synthesised in the
19
root forms a chelate with iron prior to its transport into
the soybean plant. Synthetic chelates are widely used as
carriers of iron in plant nutrition (81) to combat iron
chlorosis, a common deficiency condition in various soils
of the world. The original work of Stewart, Leonard and
co-workers (82-85) has yielded spectacular results in the
treatment of iron chlorosis through the use of EDTA and
HEDTA. The F«(m) chelate of DTPA has been found to have
similar advantages in its application as an iron carrier in
plant nutrition (86, 87). Another complexing agent which
seems to offer promise as an iron carrier is EHPG- (88).
The use of zinc chelate of EDTA in maintaining healthy
growth of plants has been reported by Perkins (63).
Excellent growth response was observed with roses by
Gerhard (89) in soil applications of Mn-ELTA chelate. The
application of chelated bas been shown by Beattie
(90) to produce an increase in both the yield and the sugar
content of grapes.
Chelates have been used to alleviate the toxicity
of heavy metal poisoning. A notable example in this field
was the early use of BAL to control the effect of arsenical
war gas poisoning (91). I*1 addition, BAL has been success
fully used for mercury and bismuth poisoning. Kety (92)
attempted to treat lead poisoning with sodium citrate,
20
since it was known that lead formed a fairly strong complex
with citrate. While citrate has been used in an attempt to
accelerate the excretion of other metal ions (93)* results
hare been disappointing since it is rapidly metabolized and
does not bind most ions with sufficient avidity. The poly
amino acids in general are extremely effective chelating
agents (9*0. Foreman and co-workers (95) have demonstrated
a tenfold increase in urinary iron excretion subsequent to
the intraperitoneal injection of calcium EDTA in rats.
This represented the first removal of iron from the mamma
lian organism by artificial means.
In recent years considerable attention has been
given to desferrioxamine, a metabolite of a species of
atreptomyces which specifically binds with ^( u j ) "to form
ferrioxamine (96). Heilmeyer and Wohler (97) treated hemo-
chromatotic patients with desferrioxamine. The ohelated
iron is excreted by the way of the kidneys because of its
relatively low molecular weight. The drug inhibits the
intestinal absorption of iron but does not appear to affect
any other trace metal•
The chelating agent BAETA has a favorable ratio for
binding of calcium and strontium than other known chelating
agents. The logarithm of the stability constant of the
chelate for calcium is 10.06 and for strontium is 9*36 as
compared to those of 10.70 and 8.63 for the EDTA chelates
21
of calcium and strontium. Spencer, et al. (98) studied the
effectiveness of the oalclum salt of BAETA in enhancing
radio strontium excretion. Intravenous calcium gluconate,
as well as Ca-BAETA was ineffective in the hypercalciuric
patient, while both agents were equally effective in a
patient with normal calcium metabolism.
While the therapeutic use of chelates in metal
poisoning is quite successful, the efficacy of chelates in
enhancing mammalian absorption of iron through the intes
tinal wall is not as obvious. Larson (99) has shown that
the addition of a 15:1 molar ratio of the disodlum salt of
EDTA to iron in the food of rats decreased the net absorp
tion of daily iron and increased the excretion of iron in
the urine, normally almost nil. However, Mills (100, 101)
has noted that organic complexes of copper present in
herbage are better absorbed in the rat than the non-chelated
ion. It is evident that the efficacy of a particular che
late is dependent on its chemical nature, as well as the
organism concerned.
CHAPTER II
OUTLINE OP PROBLEM AND METHOD OP ATTACK
While the literature is replete with clinical
studies on iron metabolism (11, 102), little is known of
the mechanisms by which iron is moved from the lumen of
the gut into the blood and then to the various cellular
compartments and binding entities within the body. There
is a lack of agreement among the investigators in this
field and the theories of these mechanisms are varied. We
decided to attack the problem of transmucosal transport by
studying the in v-ivg and in vitro effect of transferrin on
the flux rate, possible participation of an endogenous
chelate, and the dependence of the rate of transport and
distribution on the stability constant and conformation of
the iron-chelates•
Transport of trace metal ions, including Fe^+,
2+ 2+
Cu , and Mn is not intimately linked to respiratory
energy (103). Their transport is facilitated by soluble
chelates of low molecular weight (33). In the search for
sequestering agents of iron which would be useful in the
elucidation of the role of chelates in the transport of
trace metals, it was found that fructose and other sugars
22
23
can form highly stable complexes* These are rapidly
absorbed when presented to the mucosal wall (10).
There Is a general lack of information as to the
chemistry of these complexes. A systematic chemical Inves
tigation was undertaken to gain Insight Into their unique
chemical properties which make them particularly suitable
for their role in the biological systems.
Both in vivo and in vitro experiments were under
taken to study the regulation of iron transport while
chemical studies were pursued in the areas of chemloal and
physicochemical properties of iron-sugar chelates. Pour
principal aspects have been presented in this dissertation:
A. Regulation of iron absorption and utilization
B. Chemical characterization of iron-sugar
complexes
C. Stability constant of iron^jjj)-fructose at
acid pH
D. Partial periodation of iron-sugar complexes
The techniques utilized and the results obtained
are presented in the following chapters.
CHAPTER III
EXPERIMENTAL AND RESULTS
Regulation of Iron Absorption and Utilization
Respite many studies of Iron absorption, little is
known about the exact process of iron uptake by the intes
tinal mucosa and its transfer from the gut lumen across the
mucosal cell to the vascular supply of the Intestine. Most
of the studies in the past have been concerned with the
factors promoting or retarding the uptake of iron rather
than the nature of the process itself. Furthermore, the
concept of "mucosal block," with the assignment of ferritin
as the chemical cornerstone for this theory, has largely
dominated thought about the mechanism of iron absorption.
Schacter, et al. (8) demonstrated that everted seg
ments of the small intestine of the rat are able to trans-
59
port Fe from the mucosal to the serosal surface against
concentration gradients when incubated in vitro. They
stated this process as active transport dependent upon oxi
dative metabolism, and, they implied that it required the
generation of phosphate bond energy. These experiments
could not be confirmed by other workers (35).
Saltman, et al. (10) proposed a control mechanism
2M-
25
Involving binding of iron to transferrin subsequent to its
mucosal absorption as chelate bound complex. Brown (35)
et al. have been unable to demonstrate any transferrin
effect on iron absorption. In order to clarify these
results and further test the hypothesis of Saltman and
co-workers, in vitro and in vivo experiments using various
chelating agents have been performed.
In Vitr»
In vitro transport of iron was measured utilizing
the apparatus shown in Figure 2. Bats were starved twenty-
four hours prior to decapitation and exterpation of the
duodenum, approximately 12 centimeters long. This section
was everted and washed several times in 37°C saline. It
was then mounted as shown in the everted position, placed
in the apparatus, and the saline solution added. Stirring
was accomplished by means of bubbling air as shown. The
total elapsed time from killing to addition of isotopes was
approximately fifteen minutes. The incubating medium on
both sides of the intestinal loop was standard saline,
0.9 per cent sodium chloride. The entire apparatus was
submerged in a constant temperature bath at 37°C. The
59
Fe -chelates were added to the mucosal side in a 2.0 ml
6
volume. The total activity was approximately 1.1 x 10
counts per minutes. At indicated time intervals 1.0 ml
aliquots were drawn from the serosal side and radioactivity
26
serosal medium
luminal medium
air
intestinal segment
Figure 2.— Apparatus to study in vitro
transport of iron.
27
determined using a Nuclear Chicago scintillation well coun-
l?Q
ter with pulse height analyzer set for the 1.016 mev Fe
peak. The aliquot was then returned to the serosal com**
partment and the reaction continued. At the end of the two-
hour period of incubation *f0 mg apotransferrin, obtained
from Behring Werke Industries in Germany, was added to the
serosal side in 2.0 ml buffer solution. The incubation was
continued for another hour and 1 ml aliquots taken for
determination of activity. At the conclusion of the experi
ment the intestinal section was removed and washed with 0.1
N HC1 to remove any iron adsorbed on the surface. The
entire segment was blotted dry between pieces of filter
paper, weighed, and Its content of radioactivity determined.
In Vivo Experiment
in TtTQ experiments were carried out as follows*
Bats weighing approximately **00 grams were starved twenty-
four hours prior to the experimental procedure. They were
anesthetized by intraperitoneal Injection of OA ml Nembutal
(60 mg per ml). The abdominal fur was removed, the abdomen
opened, and a 12 cm section of proximal small intestine was
tied off by means of surgical thread. No washing on this
segment was performed. Ferric chelates were prepared by
adding lOOyug FeCNO^)^ and 100 fold excess of chelating
agents In 2 ml water containing Fe^ (1.0 x 10^ counts per
minute of total radioactivity). This solution was injected
28
Into the segment. The attached intestinal loop was
returned to the inside of the abdomen, a surgical sponge
containing warm saline was placed over the incision, and
the animal kept warm for the two-hour period of transport.
At the conclusion of the experiment, the entire section of
the Intestine containing the radioaotivity was excised by
cutting on both sides of the loop.
This segment was held over a 100 ml beaker, one of
the ends cut off, and the contents drained into the beaker.
The exterior of the loop was carefully washed with warm
saline, the loop everted and washed again in dilute acid.
The entire loop was then blotted dry, weighed, and radio
activity determined. The radioactivity in the solution
initially present in the loop was determined on a suitable
aliquot following dilution. A 2.0 ml aliquot of blood was
removed from the inferior vena cava using a syringe with a
#20 needle. The liver, spleen and kidney were exterpated,
washed and radioactivity determined. The contents of the
bladder were removed as completely as possible with a
syringe and radioactivity determined. Radioactivity
remaining in the carcass after removing these organs could
not be determined accurately. However, the value was very
small as measured by a hand monitor system.
When transferrin was added to the in vivo system,
0.7 ml transferrin solution (20 mg per ml in saline) was
29
Injected directly Into the Inferior vena cava. To ascer
tain whether or not the Iron binding capacity of the cir
culating blood was changed It was directly determined by
the method of Peters, et al. (104) on a separate series of
animals.
The percentage of activity transported was obtained
by summing up the activity of each of the organs measured
plus the total content of the bladder. The amount of Fe in
the gut was not Included In the calculation.
Results
The rate of transport of iron In the In vitro sys
tem is a function of the particular chelate utilized as
shown in Figure 3* The amount of iron Is given as the
percentage of total dose transported per gram of Intestinal
loop injected. Each point represents the average value for
the number of experiments indicated in the parenthesis
below the name of the chelate. At 120 minutes, indicated
by the arrow, transferrin was added to the serosal side.
It would seem that no significant change of rate of trans
port was observed.
The rate of iron transport as a function of a con
centration of iron citrate complex is shown in Figure 4.
Since the amount of radioactivity was the same in each
case, the amount of iron transported per unit time is
directly proportional to the concentration of iron over the
30
EDTA (8)
NTA (*f)
Citrate (5)
Ascorbate (6)
Pe2+ (2)
Figure 3.— Percentage of the total iron trans
ported per gram of tissue for the various chelates tested.
The number of experiments is shown in parentheses. In all
experiments apotransferrin was added to the serosal com
partment at 120 minutes. All data represent in vitro
experiments•
o
n
w
s
C O
E-i
30
20
10
180 120
MINUTES
% TRANSPORTED / GRAM
31
/Ug Fe
50
30
10
io3 .
io2 & 10
10
100 200
MINUTES
300
Figure b.— Rate of iron transport in vitro as a
function of conoentration of iron citrate.
32
range of 10 /*g to 10,000 y«g of iron. If there were a
facilitating "carrier,1 ' it could not be saturated. Simple
diffusion appears operative in this reaction. The amount
of iron transported in vivo during a two-hour time interval
as a function of chelate administered is presented In
Table 1• Also shown in this table is a comparison of the
in vitro and in vivo transport rates.
The in vivo distribution of the isotopes in various
organs as function of the chelate following two hours of
Incubation are shown in Table 2. It should be noted that
there are significant differences, not only in the rate at
which each chelate moves, but also the control of the
ultimate sites of iron deposition in the organism. Par
ticularly noteworthy is the large fraction of EDTA-iron,
which was immediately secreted into the urine and lost to
the animal. On the other hand, the NTA seems to be
extremely well utilized as is indicated by the concentra
tion in the blood and Its absence from the urine.
One of the chelates, NTA-Iron, was used in vivo
to study the effect of added transferrin to the periferal
circulation. In Table 3 the distribution for these che
lates when transferrin is present or absent is seen. The
iron binding capacity of the control plasma was l6l ^ug per
100 ml of blood, while that in the transferrin-injected
animal was 3**-9 /*g.
TABLE 1
PERCENTAGE OP THE TOTAL IRON TRANSPORTED IN
TWO HOURS PER GRAM OP WET TISSUE FOR
THE VARIOUS CHELATES TESTED
NTA EDTA CITRATE AMMONIUM
CITRATE
ASCORBATE FRUCTOSE FeSO^
in vivo 22.6
17.*f
10.1 .7** .89
.78
in vitro 17.1
20.1
11.5 9.31
6.98
C M
C M
•
.7M-
u>
u>
TABLE 2
THE DISTRIBUTION OR IRON IN VARIOUS ORGANS AFTER IN VIVO INCUBATION
THE VALUES REPRESENT PERCENTAGE OF THE TO*Ai
TRANSPORTED IN THESE ORGANS ONLY
NTA EDTA CITRATE AMMONIUM
CITRATE
ASCORBATE FRUCTOSE FeSOi*
Liver
35.9 5.1 17.7 18.7
13.3
15.0
12.3
Spleen 0.8 0> 1.8
3.1
2.i f 5.0 2.1
Blood 57.0 37.i f 76.7 7*f. 3
78.8 7*f .7
80.0
Kidney 7
15.8
3.*fr
3.9
if.l 5.3 if.9
Bladder 1.6
i f 1.3
oA 0.0 l.if 0.0 0.7
u>
-r
35
TABLE 3
THE DISTRIBUTION OP IRON IN VARIOUS ORGANS AFTER
IN VIVO INCUBATION USING NTA CHELATE WITH AND
WITHOUT THE PRIOR ADDITION OF APOTRANS-
FERRIN TO CIRCULATING BLOOD
Site of
Deposition
Apotransferrin
+
Liver 11.5 35.8
Spleen b.h .8
Kidney
6.3 h.7
Bladder 7.8
1.5
Blood 70.0 57.0
36
Endogenous Ligands
The previous experiments suggested the possibility
of an endogenous chelate participating in the transport of
iron through the mucosal wall• Experiments were designed
to assay for the presence of this chelate in rabbits.
Rabbits weighing between 2.3 and 3.3 kg were
anesthetized with 20 per cent urethane in saline by injec
tion into the marginal ear vein. The dosage used was 7.5 ml
per kg. If additional anesthetic were needed, 0.1 to
0.2 ml of nembutal (60 mg per ml) was added in addition to
the urethane. The peritoneal cavity was entered and a
20 cm section of the duodenum immediately adjacent to the
pyloric valve was located. This section was isolated from
the remainder of the gut with double ligature, making sure
that the vascular bed was left intact.
Hypodermic needles (covered with a short length of
polyethylene to prevent undue damage to the mucosa) were
Inserted in each end of the isolated segment. Using the
needles and a syringe, the lumen was washed with 25 ml por
tions of physiological saline at 37° C. Gentle pressure
eliminated excess saline from the gut after which the
distal needle (in reference to pyloric) was removed, and a
ligature tied above the point of insertion. Pour milli
litres of saline which had been rendered free from oxygen
containing 10 mg of FeSO^ mixed with Fe^ tracer were then
37
injected into the proximal end of the gut loop, which in
turn was closed with thread. The needle was left in posi
tion.
After two hours, the gut was opened, the contents
drained and centrifuged, and the following testB were
carried out with the supernatant:
1. Two drops of 10 per cent oOrt,'-dipyridyl in alco
hol were added to 2 ml of the supernatant. No
color was visible which suggested the absence
2+
of free Fe
2. The pH was reduced to b,5 by the addition of
1 N HC1 in a pH meter and then the above test
was conducted. It showed a pink color, indi-
2+
eating the presence of Fe in bound form in
the supernatant obtained from the gut.
3. Two drops of a 5 per cent thioglycolate solu
tion were added to 2 ml of the supernatant
before adding o^pC'-dipyridyl and was compared
with 2. No appreciable change of color was
visible which indicated the absence of Fe^+.
b. Two drops of K^FeCCN)^ solution were added to
2 ml of acidified supernatant. A bluish color
was visible indicating the presence of com-
plexed Fe2+ in the gut content.
As this suggested the presence of Fe^+ ion in
38
'bound form in the supernatant, purification of this com
pound was attempted.
The supernatant was poured in a Visklng sac and
dialyzed against ion-free water for twenty-four hours.
Anerbio conditions were maintained by bubbling Ng through
the external solution. The dialysate was lyophylized and
the product obtained as a dry powder containing much salt.
This was purified by ion exchange chromatography using
Amberlite ISA V01 ion exchange resin supplied by Fisher
Scientific Company and the elution of the endogenous com
plex was followed by the presence of radioactivity.
The distribution of radioactivity during fractiona
tion of the gut contents is presented in Table If. Ratio of
the high molecular weight to low molecular weight compound
obtained from the supernatant was 1:7*
The high voltage paper electrophoresis of the puri
fied dialysate was run in a Servonuclear Corp. High Voltage
Electrophoresis Apparatus using ET *f0 Electrophoresis Tank
and 5 Kv power supply for twenty-five minutes at !f,500
volts in NaHCO^/CC^ buffer at pH 8.5 on Whatman no 1 paper.
The paper was dried and the radioactive spots were located
by radioautogr&phy using Kodak single-ooated medical X-ray
film (blue sensitive). The radioactivity moved as a single
distinct separate entity. No other radioactive fraction
was found. There was not enough of this compound to carry
39
TABLE b
DISTRIBUTION OP RADIOACTIVITY IN THE
CONTENTS OP THE INTESTINAL SEGMENT
Total radioactivity placed
in the intestine
131*990 c.p.m.
Contents of segment
centrifuge
20 min. at 12,000 x g
Precipitate
Hydrolyzed
with 6 N HC1
for 10 min*
Supernatant
107,^0 c.p.m.
Dialyzed
2*f hrs.
Precipitate
**,500 c.p.m.
Solution
IS,100 c.p.m.
Dialyzed Sol. Dialysate
2 ml 25 ml
High mol. wt. comp. Low mol. wt. comp.
13,*+00 c.p.m. 9^*000 c.p.m.
ko
out further experiments.
The Relative Effect of Chelatee on
Iron Uptake by Liver Slices
Saltman, et al. (33) had measured the uptake of
iron by rat liver slices from Krebs-Ringer phosphate buffer
solution. The iron in this solution was added as ferric
ammonium citrate. This compound is a charged molecule
whose structure has not been elucidated fully. It is proba
ble that the iron is bound to the citric acid as a negative
complex (81) with the positively charged ammonium groups
providing electrical neutrality.
If the relative rate of uptake of iron by liver
slices is a function of the size, charge and stability con
stant of the iron chelate presented, the kinetics of accumu
lation would be considerably different if the iron were
offered as the small chelate molecule rather than as the
non-diffusible transferrin molecule. The great affinity of
transferrin for iron would sequester a considerable amount
of the metal in the medium surrounding the slice.
To test the above hypothesis, the rate of uptake of
59
Pe ' by liver slices was followed when the metal is given
59
as Pe ^-transferrin in presence of low molecular weight
59
ligands and as Pe -chelate of low molecular weight in
absence of transferrin.
A 2.5 kg rabbit of unknown breed was sacrificed by
hi
decapitation. The liver was perfused in situ with 100 ml
of 0.9 per cent NaCl solution via the portal vein. It was
then excised and rapidly homogenized. The homogenate was
finally dialyzed against ion free water for twenty-four
hours. The dialysate was dried under reduced pressure
which gave a white residue weighing 206 mg.
One hundred and sixty millilitres of human serum
was centrifuged to remove the red cells. One hundred
59
milligrams of freshly prepared Fe -citrate was added to
the clear serum. The serum was dialyzed for twenty-four
hours against a continuous flow of distilled water at 5° C
to remove excess citrate. After dialysis 1 ml of serum
contained 60,680 c.p.m.
The serum was equally divided into four parts.
Liver dialysate (206 mg), fructose (127 mg), citrate
(305 mg) and sodium chloride (*+2 mg) were separately dis
solved in these four parts. The pH was adjusted to 7.b in
each.
Perfused liver was excised and rapidly sliced using
a conventional Stadie-Riggs microtome. The slices were
blotted dry with filter paper and weighed. Slices weighing
between 60 and 100 mg wet weight were incubated in 3 ml of
serum mixture. At various intervals, 15 min., 30 min.,
1 hr. and 2 hrs., the slices were taken out, washed twice
with saline and once with distilled water and placed in
b2
counting vials for determining the radioactivity incor
porated.
A set of four beakers containing liver dialysate,
fructose, citrate and NaCl was used to record the zero time
adsorption in which the slices were ;}ust immersed and
immediately taken out and washed as previously.
The data are presented in Figure 5. Highest rela-
59
tive Fe incorporation was observed with citrate. Incor
poration in the presence of fructose was comparable to
control sodium chloride.
59
The relative rate of Fe uptake by the chelates in
the absence of transferrin in the medium was done as
follows*
If the chelate facilitates transfer to cells from
59
transferrin then, Fe ^-chelate when given alone should be
incorporated in the liver slices. The following experiment
was designed to test this hypothesis.
59
Fe 7-complexes were made with 1.5 mg Fe(NO^)^ dis
solved in 3 ml physiological saline with 150 mg each of
liver dialysate (70.5 x 10^ c.p.m.), citrate (60.6 x 10^
c.p.m.), fructose (70.2 x 10^ c.p.m.) and control sodium
chloride (66.3 x 10^ c.p.m.). Liver slices were incubated
in the respective solutions at pH 7.*+ for the indicated time
intervals. The slices were removed, washed twice with
saline and once with distilled water and radioactivity
^3
3600
2700
o
o
rH
^ 1800
B
m
900
O —
o
a
6o 30 90 120
MINUTES
Figure 5.— Kinetics of relative^incorporation of
radioactivity into liver slices from Fe59-transferrin
(-o- Citrate, - a - liver dialysate, -x- fructose,
-o- control NaCl).
¥t
determined. Another set of four beakers containing the
59
Fe -complexes vas used to record the zero time adsorption
in which the slices were immersed, immediately taken out
and washed as previously. The data are presented in
Figure 6. Iron as ferric fructose showed the highest rela
tive incorporation into the liver slices when presented
alone.
Chemical Characterization of
Iron-Sugar Complexes
Sugfljr~Me ta.1 Complexes
Chelation of the metal plays a fundamental role in
its availability for transport across the intestinal mem
brane. Charley, et al. (10) have shown both Fe++ and Fe+++
when suitably chelated with fructose are rapidly absorbed
when presented to the mucosal wall.
There is a general lack of knowledge about the
chemistry of these sugar chelates. Sugar complexes have
been previously reported. These have been either poorly
soluble (105“107), or high molecular weight colloidal
preparations (108-110) of uncertain structure. One with
practical significance is obtained when calcium hydroxide
is added to 5 ~ 10 per cent aqueous solutions of fructose
or invert sugar. In this way Peligot (111) and Winter
(112) obtained fine needles of the relatively insoluble
fructose compound C^H^gO^-CaO^HgO, known as calcium
1*5
3000
2250
o
o
H
1500
v
a
p «
o
750
o-
o
60 30
90 120
MINUTES
Figure 6.— Kinetics of relative incorporation of
radioactivity from Fe?9 chelate Fe ^-fructose,
-A,- Fe-^- liver dialysate, -o- Fe'-'-Citrate, -o- Control
Pe59 NaCl).
*t6
f rue to sate or levulate. It is unstable on drying and is
readily decomposed by carbon dioxide. The preparation of
this calcium compound has been improved by Jackson, Silsbee
and Prof fit (113) and affords a convenient way of purifying
fructose, in particular for separating it from glucose,
whose calolum compound is relatively soluble. Mills (llV)
has demonstrated the formation of soluble calcium complexes
with variety of polyols. Recently Charley and Saltman
(115) have shown the chelation of calcium by lactose and
indicated its role in transport mechanisms.
A systematic chemical study of ferric fructose was
undertaken.
Spectral Evidence for
flnelate floraatfon
Dilute PeCNO^)^ solution is a pale yellow at acid
pH. Upon the addition of excess sugar, the color darkens
considerably and there is a slight decrease in pH. Such a
change in spectral properties has been used to establish
the presence of a chelate (70). It was discovered that
equimolar proportions of iron to monosaccharide were com
pletely ineffective in preventing hydroxide precipitation
under alkaline conditions, although the spectral change was
apparent. The color of the ferric fructose is dark yellow
under acid conditions. The hydrated ferric ion absorption
is very strong in the shorter wave lengths so that the
b7
presence of the darker complex which is apparent visually,
can not easily be demonstrated with the spectrophoto
meter. Same spectral properties of this ferric fructose
chelate can be observed by measuring the difference spec
trum between a 10 ^M FeCNO^)^ solution at pH 2.5 and the
same concentration of FeCNO^)^ in the presence of 16-fold
molar excess of fructose at the same pH (Figure 7). The
increased absorption of the complex at 370 m^ Is clearly
demonstrable, despite a close similarity of the spectra for
the individual solutions.
Tjtrimetric and Retentionstrie
Evidence for Complex Formation
The displacement of one or more hydrogen ions from
the sugar molecule during the formation of a fructose che
late was directly demonstrated by the titration curves
shown in Figure 8. Fifty millilitres of a solution 0.25 M
fructose, 0.01 N HC1 and 0.012 M in either FeCl^ or FeClg
were titrated with 0.11 N NaOH. In order to be certain that
2+ 3+
no interaction of Fe and Fe would interfere with the
titration, 50 ml of a third solution 0.25 M fructose,
0.01 N HC1, 0.006 M Fe^+ and 0.006 M Fe2* were likewise
titrated. The pH was determined with a Radiometer Model-22
pH meter. The titration curves reveal that 3 H+ are
directly displaced from the sugar and ^0 for each Fe^+
initially present. Complex formation is complete for Fe^+
% transmittance
0
20
wo
60
80
100
550 500 W50 Woo 35b
Wave length (m )
Figure 7.— Difference spectrum of ferric fructose
compared to ferric ion. All solutions were adjusted to
pH 2.5
Ferric fructose
rency spe9tru.n1
*t9
10
8
6
2
0
T77T U72
mequiv NaOH
A 2 ^
Figure 8.— Titration curves of Fe + and Fe
separately and together in the presence of excess fructose.
The initial volume in each case was 50 ml of a solution
0,25 M fructose, 0,01 N HC1, and (a) 0,012 M Fe2*»
(b) 0.006 M Fe and 0.006 M FeJ+, or (c) 0.012 M Fe^ .
50
by pH 3.5. In the study with Fe^+, 2H+ are displaced per
iron. However the complex formation does not begin until
about pH 7.5 and is essentially complete by pH 8.5. At all
times the solutions were clear. No precipitate or colloid
formation could be detected. The titration curve for Fe^+
2+
and Fe in the same solution indicates that there is no
interference between oxidized and reduced iron ions.
Potential measurements over a wide range of pH were
• 2+
made in a solution containing 0.006 M Fe and 0.006 M Pe .
The pH and redox potential were measured simultaneously
using two Model-22 Radiometers, one fitted with a glass
electrode, the other with a platinum electrode, using a
standard calomel half cell. The initial pH of the solution
was 2.0. It was rapidly titrated with 0.11 N NaOH. The
values obtained are presented in Figure 9. However, when
the same measurements are made in the presence of 0.25 M
3+ 2+
fructose it can be seen in Figure 9» that the Pe and Pe
are in equilibrium over a wide range which extends to pH
10*5. The break in the slope of the line at pH 8.25 ie
2+
caused by the formation of the Pe complex with the con
comitant lowering of the concentration of this free ion in
solution. The slope of the redox potential vs pH line of
Figure 9 is that which one would find if 3 H+ were
liberated in the chelation of Fe^+ by fructose and con
firms the tltrimetric evidence presented above. This
51
600
300
o
Control
-300
Fructose
-600
b 2 10
Figure 9*— Redox potential as a function of pH.with
and without fructose. Initial solution was 0.006 M Fe^*
and 0.006 Fe2*; fructose, when present, 0.25 M.
52
system is not freely reversible and hysteresis effects were
observed. Therefore only qualitative conclusions can be
drawn.
Isolation and Test of Purity
of ferric Fructose
The ferric fructose complex can be precipitated
from aqueous solution by the addition of ethanol to a final
concentration of 80 per cent (v/V). The complex was pre
pared by rapidly titrating 20.0 ml of a solution of 0.1 M
in FeCNO^)^ and 1.6 M in fructose to pH 8.5. Eighty milli
litres absolute alcohol were added, the pH checked again,
and the precipitate collected by centrifugation. The pre
cipitate was redissolved in a small volume of water, the
pH adjusted to 8.5* and this clear solution again made 80
per cent in ethanol. The precipitate formed was collected
by centrifugation, washed successively with ethanol, ace
tone and ether and dried over ^Otj.
Ferric fructose thus obtained was chromatographed
on glass fibre paper using 70 per cent ethanol (v/v) for
development. It was sprayed with 5 per cent thiocyanate
solution and resorcinol (137) for Fe and fructose respec
tively. Both Fe^+ and fructose were located in the identi
cal place. No free fructose or iron spots were seen on the
paper. When ferric fructose was brought to pH 1 by HC1 and
slowly raised to pH 7 by the addition of NaOH, iron
53
precipitated as its hydroxide all sugar was recovered as
fructose analyzed by Roe's method (116).
Dialysis of ferric Fructose Chelate
59
Pe * was used to study the rate of ferric fructose
dialysis through a Visking sac compared with that of Pe-
EDTA at equal concentrations. Perric fructose was prepared
by mixing 1 ml 10 2 M Pe(R0,)^ containing 2.78 x 10^ c.p.m.
59 -p
Pe and 1.6 x 10 M fructose in 3 “1 tris buffer at pH
7.2. Pe-EITA was prepared by mixing 1 ml 10~ * 2 M FeCNO^)^
containing 3*07 x 10^ c.p.m. Fe^ and 1 ml 10 2M EDTA in
3 ml 10~2 M tris buffer at pH 7.2. One ml of the above
solution was placed in a Visking sac and dialyzed against
-2
25 ml of 10 M tris buffer with constant agitation.
Dialysis of ferric fructose also was carried out against
buffer containing 0.25 M fructose outside the sac. At
time intervals an aliquot of 1 ml was taken out from the
outside and radioactivity determined in a Nuclear Chicago
scintillation well counter using a pulse height analyzer
set for the 1.098 Mev. peak.
The data are presented in Figure 10. Fe-EDTA
reached 97 per cent equilibrium in thirty hours. Ferric
fructose with 0.25 M fructose outside diffused slowly and
reached the equilibrium comparable to Fe-EDTA in 180 hours.
Perric fructose with no fructose outside diffused very
slowly and after 2^-0 hours only 5 per cent of equilibrium
5k
h
X>
100
80
60
20
30 60 90 120
HOURS
150 180 210
Figure 10.— Percentage equilibrium reached as a
function of time in0the dialysis of Fe-chelates. -x-
Fe-EDTA against 10“2 M tris buffer, -O- Ferric fructose
against 10“2 M tris buffer containing 0.25 M fructose in
the outside, -A- Ferric fructose against 10”2 M tris
buffer.
55
was established. Time required to reach 50 per cent equi-
I
lihrium Ct^/2) was ten hours for Fe-EDTA and that for ferric
fructose with 0.25 M fructose outside was seventy hours.
Elemental Analysis and
Molecular Weight
Elemental analysis for C, H, Fe, Na and 0 by differ
ence, was performed (Microanalytisces Laboratorium in Max
Planck Institute fur Kohlenforschung, Mulheim, and Labora
torium Fresenius, Wiesbaden, Germany) after thorough drying
in high vacuum. Following is the analysis of ferric fruc
tose given as percentage of total weight:
Ferric Fructose C H O Fe Na HgO
Experimental 26.1 3.6 *f3.5 22.9 ^.0 10.5
Theoretical value based on
the empirical formula 27.5 ^.2 ^2.8 21.*+ *+.2 12.1
(C6H1006)2re2(H20>3(OH)3H8
Theoretical value based on
the empirical formula 28.6 M+.5 22.3 --- 12.5
( C6H1006 ) gFte 2 (H20 > 1 * ( OH) 2
The sample of ferric fructose used for the elemental analy
sis was isolated free from fructose at pH 9.0 and dried
over P2O5. H2O mentioned in the table was removed by drying
in high vacuum. The oxygen value given does not include
that from water of hydration.
56
Since ferric fructose is insoluble in the organic
solvents commonly used for molecular weight determinations
by freezing point lowering, an aqueous solution containing
50.3 mg/si of the material used in the analysis was pre
pared and its freezing point lowering measured as 0.3l5°«
If one assumes complete dissociation of the Na-salt of the
complex, the experimental value for the molecular weight is
59^. However, preliminary molecular weight determination
of ferric fructose by ultracentrifugation (Spinco Division
of Beckman Instruments) gave a value of about 7.0 x 10
which indicates a polymeric form of the complex isolated at
that pH. The presence of sodium bicarbonate as an impurity
in the sample of ferric fructose can seriously affect the
measurement of molecular weight by freezing point depres
sion method. In our experimental conditions sodium bicar
bonate could be a major single source of contamination.
Infra-rad Spectra of
Ferric Fructose
Infra-red spectra were determined to see if there
were any particular change in the hydroxyl region because
of the suspected iron binding with fructose hydroxyl
groups. It might also be possible to observe if the car
bonyl group of straight chain of fructose was involved in
chelation.
Infra-red spectra for both fructose and ferric
57
fructose were recorded on a Beckman IB spectrophotometer by
Spineo Division of Beckman Instruments. The sample was
dried before using. Five milligrams of it were mixed in
Nu^ol Mull and air was used as the reference.
Infra-red spectra of both are presented in Figure U.
There is a considerable difference between fructose and
ferric fructose in the region 3600 cm”1 to 3000 cm-1.
Ferric fructose shows an absorption peak at 1625 cm 1 which
is attributed to co-ordinated water (117). Much fine struc
ture in the spectra is seen in fructose in the region of
1500 cm"1 to 700 cm"1. Fujita (118) has examined the infra
red spectra of typical aquo complexes and they have shown a
band characteristic of co-ordinated water near 795 cm
Ferric fructose shows a characteristic band in the same
region. Ring oxygen in assymetrical C-O-C stretching vibra
tion of cyclic structures gives a strong band close to
1060 cm 1 region. Both fructose and ferric fructose show
this band.
The information obtained from the infra-red spectra
was not helpful in any further interpretation. It is diffi
cult to compare two compounds whose physical state is not
similar. Crystalline fructose shows many additional absorp
tion peaks which are nonexistent in ferric fructose in the
region of 600 cm”1 to 1000 cm”1. Infra-red spectra of many
more chelates whose structures have already been elucidated
58
Fructose
100
75
25
000 3600 3200 2*f00 1600 l*f00 1000 800 ifOO
Wave number (cm’^)
Ferric fructose
100
75
50
25
ifOOO 3600 3200 2*f00 1600 l*f00 1000 800 *f00
Wave number (cm
Figure 11.— Infra-red spectra of Ferric fructose
and Fructose.
59
should be looked at to get a better understanding for its
interpretation.
Investigation ofEnediol Configura
tion for ferric Fructose
If the ligand between iron and fructose were
through the 1 and 2 position of fructose in an enol con
figuration, then on acid hydrolysis it should yield a mix
ture of fructose and glucose as follows!
H
i
H-C-OH H-C-OH H - C = 0
i ti i
C-0 C-OH H-G-OH
» i i
R R R
keto enediol aldo
A complex of iron with glucose should be identical and
react in the same fashion as the fructose chelates. The
following experiment was performed:
Perric fructose, 0.01 gm, was dissolved in 5 ml
water and the pH adjusted to 1.8. One and one half milli
litres 0.1 M EDTA was added to it. The pH was then adjusted
to 7.5 and enough absolute alcohol added to obtain a pre
cipitate. The mixture was centrifuged and the supernatant
taken to dryness in a flash evaporator. The dried material
was again mixed with alcohol and evaporated in the same
manner as before. It was again mixed with 80 per cent
ethanol and the mixture filtered. This filtrate was used
6o
for chromatography.
The solvent system was 100 ml of 80 per cent phenol
mixed with 0,5 ml of concentrated NH^OH. Aliquots of both
free glucose and fructose were cochromatographed on Whatman
Paper No• 1 along with the hydrolysate• Following develop
ment in solvent, the paper was dried, sprayed with aniline
hydrogen pthalate to detect sugars (119)* All of the
hydrolysate moved as fructose. No glucose spot was visi
ble. Conversely, with glucose-iron, only glucose could be
detected. Thus the speculation of having a chelate struc
ture with 1, 2-ene of configuration is unlikely.
Preparation of Ferric Glucose
The ferric glucose complex can be precipitated frcm
aqueous solution by the addition of ethanol to a final con
centration of 80 per cent (v/v)• The complex was prepared
by rapidly titrating 20.0 ml of a solution of 0.1 M in
Fe(NO^)^ and 3*2 M in glucose to pH 8.5 by X N NaOH.
Eighty millilitres of absolute ethanol were added, the pH
was checked again, and the precipitate which formed was
collected by centrifugation. The precipitate was redis
solved in a small volume of water, the pH adjusted to 8.5,
and this clear solution again made 80 per cent in ethanol.
The precipitate was collected by centrifugation, washed
with successive volumes of ethanol, acetone and ether and
dried over P^O-.
2 5
61
Perric glucose thus obtained was chromatographed on
glass fibre paper using 70 per cent ethanol (v/V). It was
sprayed by thiocyanate and resorcinol for Fe^+ and glucose
respectively. Both Fe^+ and glucose were identified in the
same location. No free glucose or other iron spot could be
detected on the glass fibre paper.
Elemental Analysis and
Molecular Weight
Elemental analysis for C, H, Fe and Na was per
formed after thorough drying in vacuum (Microanalytisces
Laboratorium in Max Planck Institute fur Kohlenforschung,
Mulheim, and Laboratorium Fresesius, Wiesbaden, Germany).
Following is the analysis of ferric glucose:
C H Fe Na H20
Ferric glucose 26.99 ^.*+1 20.91 2.17 13.79
Oxygen value was determined as *f5.52 by difference and does
not include that from water.
Since ferric glucose is insoluble in the organic
solvents commonly used for molecular weight determination
by freezing point lowering, an aqueous solution containing
26.^ mg/ml of the material used in the analysis was pre
pared and its freezing point lowering measured as 0.18°.
If one assumes complete dissociation of Na-salt of the
complex, the experimental value for the molecular weight
is 57*+. However, the uncertainty in the molecular weight
62
determination by freezing point depression method is great
as disoussed earlier.
Acetvlation of Ferric Qluoose
in Pyridine
One method to determine the number of hydroxyl
groups of glucose involved in the chelation with iron would
be to acetylate the complex and measure the number of moles
of acetyl groups blocking the hydroxyl groups. Attempts
to acetylate glucose and ferric glucose in pyridine were
carried out as follows:
One gram of ferric glucose was added to 10 ml of
anhydrous pyridine and then 3.7 ml of acetic anhydride were
added with shaking. Perric glucose did not go into solu
tion immediately. After the initial reaction subsided, the
solution was warmed to for ten minutes under a reflux
condenser and complete solution was observed. The mixture
was cooled and poured into *f0 ml ice water. Initially a
very fine precipitate was observed. Twenty millilitres of
6 N HC1 were added and the mixture stirred for ten minutes.
A precipitate was obtained. The mixture was filtered and
the residue was washed five times with distilled water and
dissolved in warm 95 per cent ethanol. The ethanolic solu
tion was filtered warm and freed from insoluble material.
The filtrate was poured in a 2? ml flask and stored at *+°C
to permit crystallization.
63
Crystals were obtained from both glucose and ferric
glucose, m.p. 130° - 131°C. It would appear that both com
pounds yield the same acetyl derivatives. In this case it
was confirmed as pentaacetate by studying the mixed melting
point with the authentic sample. During the acetylation
reaction, the bonds between iron and glucose appear to be
destroyed, yielding the same product as that from glucose
on acetylation.
Toaviation of Ferric Glucose
The unimolar acylation of carbohydrate derivatives
containing both primary and secondary hydroxyl groups leads
to the preferential esterification of the former (120),
Among the secondary hydroxyl groups that may be present,
the one adjacent to the carbonyl group, i.e., position two
in the aldose sugars, is usually the more reactive. In
certain derivatives, the reactivity of the alpha secondary
hydroxyl group is enhanced by the nature of the substitution
to equal that of the primary hydroxyl group.
Many people have used the above method to elucidate
the structure of the carbohydrate components. Since the
acetylation reaction breaks the iron-sugar bonds, another
reagent was sought for specific hydroxyl group esterifica
tion. Compton (121) has shown that the hydroxyl group on
the 6-position is more reactive in a tosylation reaction
where the ratio of reagent to sugar is 1 : 1. In an attempt
64
to determine if position 6 is involved in chelation, ferric
glucose vas tosylated as follows:
Five gram of ferric glucose in 50 ml of dry pyri
dine was added to an ice cold solution of 5*39 g of tosyl
chloride with rapid stirring in $0 ml dry pyridine during
a period of fifteen minutes. After standing thirty minutes
at 0°, the mixture was removed from the ice bath and
allowed to stand for twenty-four hours at room temperature.
Fine crystals of the 6-tosyl derivative of glucose was
obtained from both glucose and ferric glucose which melted
at 132° - 133°C.
From the above experiment it would seem that the
6-position is not involved in the chelation and hence the
tosyl derivative is formed from the ferric glucose. In the
light of the acetylation reaction, any of these reagents
might attack the ligand itself and destroy it before ester-
ification. However there was no precipitate in alkaline
solution with ferric glucose indicating the presence of an
intact chelate in the tosylation mixture.
The Stability Constant of Iron^XTT\-
Fructose at Acid pH ' }
Ferric ion in the presence of fructose at pH 2.5
forms a complex which shows a distinct change of color that
can be measured spectrophotometrically. A difference spec
trum between the complex and ^®(xn) a" ^ P** 2*5 shows a
65
amm at 370 ny* which permits us to measure the dissooia-
tlon constant for the first ligand at the acid pH. The
method employed is an extension of that used to determine
the stability constant of ferric dipyridyl (122).
It was, however, necessaxy to develop several equa
tions and to apply the experimental data to them to obtain
the desired physical constant.
The generalized dissociation reaction for a complex
unit co-ordinated with n fructose is
FeFrn - c — — + elFt (1)
when FeFrn, F«(m) ^ d ** denote respectively the sum of
all hydrolytic forms of the complex, uncomplexed iron and
free fructose that are present. Thus F©(m) * (Fe+,f+) +
Fe(0H)+* + et cetera. Similarly, may be partially
dissociated as a weak acid. At a given pH, and assuming
that no hydrolytic polymers are present, we can write for
the apparent equilibrium constant
Fe/TTTN (Fe)
K = - -i ------ (2)
(FeFrn)
The total concentration of iron, (Fe)^ is given by*
(Fe)^. — + (FeFrn) (3)
and the apparent extinction coefficient is given by
66
fapp ” (je)t * ’Pi* * *>e*rn •f?*Prn (lt>
Where A Is the absorbanoy, 63^ and € are nolar
absorption coefficients (that for fructose being sero for
fractions of (Fe)^ present as *©(ui) and FeFrn reBP®c"
tively. By combining terms we see that:
If a plot of (Fr)>&*vs (Fr) is linear, then n = 1 and K and
Experimental iy*d Rftaul t
Fructose (Pfanstiehl Laboratories, Inc.) solutions
were prepared in ion-free water and stored under refrigera
tion until use. They were equilibrated at room temperature
before use. A standard solution of 0.01 M FeCNO^)^ • 9 ^0
(Baker & Adamson) was prepared in 0.1 N HNO^. The pH of
the wave length in question), and * M (•-•) and 5 x 1(T^ M (Fe)^ (o-o).
71
the plot over a wide range of fructose concentrations con
firms the implicit assumption that n * 1, and from the
slope and intercept of the best straight line, we obtain
* 1800 and K * 0A7.
The actual complexation reaction is:
H+ ♦ (FeFr)2* ^ Fe3+ + Fr (8)
Kjj for which is related to the above K of 0.^7 by the
hydrogen ion concentration and the fraction of free iron
present as Pe3+. For the latter correction, the appro-
priate hydrolysis constant for Fe at our ionic strength
of ca 3 x 10"3 was used (123)* We thus find Kj) ■ 96.
Partial Periodation of Ferric
Fructose ana Perric Glucose
The oxidation of specific linkages in sugars has
played a key role in the determination of structure.
Periodic acid, as a reagent for glycols (-CH0H-CH0H-), was
introduced by Malaprade (12**, 125). This reagent is
valuable in analysis of carbohydrates containing suscepti
ble groupings and also in determining if such groups have
been substituted.
It has been shown that the oxidation of D-mannitol,
galactitol and D-glucitol with a limited quantity of sodium
periodate (126) involves preferential attack on oC-T-glycol
groups (Barker and Bourne's nomenclature) (127). Similar
72
oxidation of erythritol showed that a -CH(OH).CH2OH group
was more readily cleaved than a -CH(OH).CH(OH) group, hut
the reverse was found for hexitols (128), Of cyclic com
pounds the threo-(oCT) was oxidized more rapidly than the
erythro isomer («CC) and for threo compounds the rate of
oxidation decreases with increasing length of the substit
uents on the glycol groups (129)•
Various mechanisms for the oxidation of glycols by
periodate and structures of the intermediate complex, com
pound or ion have been suggested. The generally accepted
theory (130) is that a cyclic ester intermediate, which
may be neutral, mono- or di-negatively charged is formed
from the glycol and E^IO^ or its dissociation products.
Keen and Symons (131) have found that ions such as H^IO^
or possibly, HglO^ are the major components of saturated
aqueous solution of sodium periodate. Buist, Bunter and
Miles (132) suggested the formation of an intermediate
with a puckered five-atom ring in which the iodine atom
is octahedral. They discussed the steriochemlstry and
electronic factors of this cyclic intermediate on the
equilibrium constant for the formation of the intermediate
and the rate constant for the decomposition to products.
None of the proposed structures has however been proved.
Hutson and Weigel (133) postulated that since oC -T
groups are more readily oxidized toy periodate than oC-C
glycol groups it would toe likely that, in a statistical
senBe, the 5-hydroxyl group is more readily available for
the formation of an intermediate involving and
than the 2-hydroxyl group for the formation of an inter
mediate involving eui& C(2)* Thus, of the two
-glycol groups, the 5,6-bond is more readily cleaved than
the 1,2-bond.
Poster (13*0 suggested that aldehydo form of
D-glucose is the principal one involved in complex forma
tion with borate and the pair of the 2- and V-hydroxyl
groups are sterically most favorable for complex formation.
Buist et al. (132) indicated a possibility of
anomalous reaction at pH 10. Certain discrepancies have
been found in the periodate oxidation in phosphate buffer
at pH 7.5 (135,136). Hutson (133) thought that similar
discrepancies might arise with a borate buffer and that
comparison between the oxidation of D-glucitol in 0.5 M
borate (pH 10) and 0.5 H phosphate buffer (pH 10) would be
valid. Oxidation of D-glucitol with 0.25 mole equivalent
of sodium periodate in the presence of borate (pH 10) or
phosphate (pH 10) yielded very different results from those
obtained with the unbuffered solution. Only c£ 10 per cent
of the D-glucitol was oxidized in the alkaline solutions
whereas the theoretical amount of 25 per cent was oxidized
7 * *
In the unbuffered solution. They indicated that the prin
cipal site of attack on the D-glucitol-borate complex is
the 5,6-bond. The pair of hydroxyl groups involved in the
formation of the complex are probably those on Bnd
coo*
Partial periodation was undertaken in our work in
an attempt to elucidate the structures of both ferric fruc
tose and ferric glucose. The reaction cannot be carried
out at acid pH because the complex is hydrolyzed at low pH.
In the alkaline range phosphate buffer could not be used
because phosphate binds with iron.
Preliminary experiments on partial periodation of
both fructose and ferric fructose were done maintaining the
pH at 7.0 by constant addition of 0.01 N NaOH. After half
an hour the pH was brought to 2.5 and then slowly raised by
addition of 0.1 N NaOH until dissociated iron precipitated
as the hydroxide. Excess sodium borohydride was added and
the reaction mixture kept overnight. It was filtered and
the filtrate was reduced to a small volume. It was chroma
tographed in 90 per cent acetone and 10 per cent water on
Whatman Paper No. 1. Sorbitol, ribitol, erythritol,
glycerol and ethylene glycol were co-chromatographed as
standards. The paper was developed by dipping it in alka
line silver nitrate (137). Both fructose and ferric fruc
tose showed the same spots. Since their concentration
75
might differ, a quantitative estimation of the degradation
product was sought.
Preliminary experiments to measure the periodation
of sugar was sought to determine if there was any effect
due to buffer. Both phosphate (10-^M) and trie (10 2M)
buffers were used in the reaction. Reaction mixture con-
-5 -5
tained 5 x 10 M glucose and 1 x 10 M sodium periodate in
10 phosphate and 10 -2m tris buffer at pH 7.5. The rate
of reaction was followed by the production of 10^ measured
at 300 ny* in a Beckman IXJ spectrophotometer with a Gilford
attachment using a 1 cm. light path. It was found that in
presence of phosphate buffer all 10^ was converted to 10
in fifteen minutes whereas with tris very little 10^ was
mm
produced from the reaction. 10^ was added in excess of
the molar concentration of tris buffer and the reaction was
completed in three minutes. Addition of excess sugar had
no effect. It seems probable that IOi^. binds with tris
buffer and inhibits periodation reaction. Thus, tris
buffer was unsuitable for two reasons. At such a high con
centration of 10^. , tris buffer cannot hold its buffer
capacity and at this concentration sugar undergoes total
periodation which is contrary to the concept of this experi
ment.
Phosphate buffer can be used with sugar but it is
unsuitable for the iron sugar complex since iron binds with
76
phosphate and the complex Is destroyed before periodate
starts reacting on the complex. A suitable method for
partial periodation in a non-buffered system was sought.
Both 0.^5 gram of glucose and 0.1065 gram of sodium
periodate were dissolved separately in 25 ml water and pH
adjusted to 9 with 1 N NaOH. One milliliter of each was
mixed together and the rate of oxidation of glucose was
studied in a XU Spectrophotometer with Gilford attachment.
Figure 15 shows the reaction is complete in fifteen minutes
during which time pH dropped only 0.5 unit. It is apparent
that the pH drop was not significant to break the complex.
For a quantitative determination of the degradation
product from the partial periodation reaction by the above
il. tU
standard method, Glucose-l-C , Glucose-6-C and
lL llf
Glucose-TJ-C were used. Chelates of Fe-glucose-l-C ,
llf Tk
Fe-glucose-6-C and Fe-glucose-U-C were prepared as
follows:
To 0.72 gram of glucose in 1.25 ml of water was
1 h
added 35/*G glucose-C . A solution containing 0.05 gram
of Fe(NO^)^ in 1.25 ml water was prepared. Both were mixed
and the pH adjusted to 8.5 by adding 0.05 ml 1 N NaOH. Ten
millilitres absolute ethanol were added to precipitate the
chelate. The chelate was centrifuged and the supernatant
was decanted. The precipitate was then repeatedly washed
with 85 per cent ethanol until no radioactivity was
77
0.3
0.2
0.1
5 10 15 20
MINUTES
Figure 15.— Rate of change of optical density due
to change in concentration of periodate on oxidation of
glucose measured at 300 ny* with an initial pH of 9.0.
recovered from the washings.
Ferrio-glucose O.lMt gram was mixed with freshly
prepared radioactive ferric-glucose in 5 ml water and the
pH adjusted to 9.0. In 5 ml water 0.09 gram glucose con
taining 5 of radioactivite labeled glucose was dissolved
and the pH adjusted to 9.0. Two separate reactions were
run in a reaction apparatus as shown in Figure 16. The
reaction flask was fitted with an inlet through which
nitrogen gas was passed at a constant rate. In 5 ml of
water 0.0213 g of sodium periodate (representing 1:5 molar
ratio of sodium periodate to sugar) was dissolved and the
pH was adjusted to 9.0. The solution was added to the
reaction flask containing 5 ml of the labeled chelate or
hexose solution. The outlet from the reaction vessel was
conneoted to a trapping flask containing 2 ml of dlmedon in
95 per cent ethanol. Any formaldehyde formed during the
reaction would be swept to the trap and be absorbed in the
dimedon solution. The reaction was continued for fifteen
minutes. At the conclusion of the reaction 0.189 gram of
sodium borohydrlde was added to the reaction vessel to
reduce the dialdehydes formed by periodation and the solu
tion left overnight in the refrigerator. The reaction
mixture was filtered and the washings from the residue
along with the filtrate collected. The solution was poured
into the same type of reaction flask as is shown in
79
■N,
Pigure 16.--Apparatus for partial periodation
/
Figure 16, and 1 N HC1 was added dropwise to bring the pH
to 2.5. Nitrogen was passed for twenty minutes through the
solution to remove formic acid. The trap in the terminal
end contained a solution of Hyamin to absorb formic acid.
The solution from the reaction vessel was poured through a
column containing 5 gram cation exchange resin AG 50-X8
hydrogen form supplied by Bio-Rad Co., (Berkeley, Calif.).
Elution was continued with ion free water until the eluate
was free from radioactivity. The eluate thus obtained was
evaporated to dryness in a flash evaporator at 30°C under
reduced pressure. Five millilitres 6 N HC1 and 20 ml
methanol were added to the residue and the solution was
refluxed for ten minutes. It was evaporated and the
residue mixed again with methanollc HC1. It was dried and
the process was repeated several times to free it from
boric acid.
Radioactivities in the dimedon and Hyamin solutions
were measured in a Packard Tricarb Liquid Scintillation
Counter in a solvent containing 70 per cent toluene, 30
per cent methanol, 5 grams PRO/litre, 100 mg POPOPAitre
mixed with 10 per cent Hyamin. The dried material free
from boric acid was dissolved in 5 ml of 50 per cent etha
nol. An aliquot from it was measured for radioactivity.
Table 5 shows recovery of radioactivity.
The samples obtained from above were spotted on a
TABLE 5
EE COVERT OP RADIOACTIVITY APTER PERIODATION OP
LABELED GLUCOSE AND PERRIC GLUCOSE
Glu-C^ Glu-C^
llf
Glu-U-C Pe-G-C^ Pe-G-C^ Pe-G-U-C1* *
Total
Initial
Counts
10^x c.p.m. 12
5.*f
12.6
5.7
9.98
^.5
7.72
3.5
6 M
2.9
>+.68
2.1
Total
in
Hyamin
lO^x c.p.m,
/aO
1.2
5.23x10*
11.8
5.32xl03
7.8
3.52xl03
15.7
7.06xl03
5.6
2.52xl03
12.1
5.>+5xl03
Total
Dimedon
0 0 0 0 0 0
Total
in
Recovery
lO^x c.p.m. 11.95 12 M 9.89
7.6 6.25
yUC 5.38 5.6 >+.1+6 3.^2 2.82 1.9 5
$ Loss OA 1.2 0.9 1.7 2.5
7.2
Whatman Paper No. 1 and co-chromatographed with sorbitol,
ribitol, erytritol, glycerol and ethylene glycol standards
in a solvent containing 90 per cent acetone and 10 per cent
water. One of the papers was dipped in alkaline silver
nitrate solution to Identify the polyols (137)• The other
was exposed to X-ray film to make a radioautograph. After
one week the radioautograph was developed and the spots
were identified. The developed film was plaoed on top of
the original paper chromatogram and the respective spots on
the paper cut out. The cut out papers were eluted with 50
per cent ethanol and the radioactivity measured in Packard
Tricarb liquid Scintillation Counter in a solvent containing
70 per cent toluene, 30 per cent methanol, 5 grams PPO/
litre, 100 mg POPOP/litre mixed with 10 per cent Hyamin.
The data for the distribution of radioactivity in different
polyols are presented in Table 6.
In most experiments periodate cleaved one fifth of
total sugar present. The possibility of multiple cleavage
is most unlikely. However, preponderance of fragment in
all cases makes any interpretation difficult. Ordinarily
one would expect more fragments than C^• Although the
percentage recovery in all the samples is good, the dis
tribution of radioactivity in different fragments does not
give us any rationale to put forth a justifiable interpre
tation. We have also explored the possibility of the
TABLE 6
DISTRIBUTION OP RADIOACTIVITY IN 100 Ml OP PINAL PERIODATE
REACTION PRODUCT OP LABELED GUJCOSE AND PERRIC GLUCOSE
100 yUl spot
C .p.m.
/xC
223,692
1.01
250,213
1.13
196,88*
0.88
123,175
0.56
129,9*6
0.59
95,3*5
o.*3
sorbitol
C *p |ID|
>uC
170,000
0.77
176,100
0.79
1*6,890
0.67
88,600
0.*
91,620
0.*1
67,2*8
0.30
ribitol
c.p.m.
103x /uC
12,*80
5.62
12,270
5.53
; ;
3,622
1.63
2,959
1.36
erythritol
c.p.m.
lO^x /»C
; ;
2,*30
1.09
2,580
1.16
:
2,33*
1.05
2,359
1.06
glycerol
c.p.m.
lO^x juC
3,750
1.69
6,550
2.95
—
3,001
1.3
2,2*+0
1.0
glycol
c.p.m. *2,500
1.91
**,000
1.98
20,*50
0.92
25,2*9
1.09
22,788
1.03
1*,289
0.6*
Extra
c.p.m.
lO^X yUC
3,196
1.**
9,*30
*.25
3,13*
l.*0
7,871
3.*
*,520
2.02
*,50*
2.0
Oo
UJ
fragments forming some form of formate ester which migrates
in the same region as the C2 fragment upon chromatography
in the acetone twater system. A different solvent system of
80 butanol* 22 acetic acid* 50 water (v/vA) (138) gave
similar results indicating C2 being the major fraction in
the partial periodation reaction besides the unreacted
components.
CHAPTER IV
DISCUSSION
Many new concepts in our understanding of mucosal
transport of iron have evolved from the studies presented
in this dissertation. The rate of transmucosal flux is a
function not only of the concentration of the complexes in
the intestinal lumen hut also of the chemical nature of the
chelates used to complex the iron. Brown and Justus (35)
using everted gut loops, presented evidence for passive
mechanism similar to that observed for iron accumulation by
liver slices (31“33). Dowdle, Schaoter and Shenker (8)
using a similar system to that of Brown and Justus claimed
that metabolic energy is required for iron accumulation.
They used fructose in their medium and addition of ascorbic
acid to the medium markedly increased the active transport
59
of Pe in vitro. However, the stimulation of transport
may be due solely to chelation of iron by ascorbic acid and
fructose and not due to any metabolic energy as indicated
by these authors.
Our findings using the in vitro gut section indicate
that a passive diffusion process is operating rather than
an active pumping mechanism or facilitated diffusion. The
85
86
iron binding capacity of the blood, as apotransferrin, has
no effect on the rate of movement of iron-complex from
lumen into blood • The level of transferrin binding capac
ity of the blood does, however, determine the fraction of
the absorbed iron which remains in the blood rather than
entering the depot sites of the liver or other organs.
This demonstrates that apotransferrln can successfully com
pete with the liver cells for available iron as would be
expected if transferrin were the intermediate in the trans
port of iron from storage depots to the sites of hemopoesls.
In comparing the availability to the tissues of
absorbed iron complex, it was found that some of the che
lates allow better utilization than others. For example,
iron-EDTA and iron-NTA complexes are both rapidly absorbed
into the blood. The iron-EDTA is excreted almost as
rapidly as it is absorbed, while the iron-NTA is largely
retained in the animal and deposited in the liver. Com
plexes of iron such as NTA should be of better therapeutic
benefit in the treatment of iron deficiency syndromes than
compounds such as EDTA, even though the rates of absorption
are comparable.
The data suggest that the relatively tight control
of iron absorption seen in normal animals is due to both
Insolubility of iron and the poor absorption of complexes
formed from dietary and endogenous ligands rather than
87
mucosal blockage of iron uptake.
Isolation of an endogenous chelate from the intes
tinal secretion suggests its participation in the transport
of iron through the mucosal wall. This chelate is a low
molecular weight compound and can bind iron. Our work is
further supported by Brown, et al. (6) who showed the
59
presence of Fe bound to serine and glycine in the mucosal
wall sorapings. However, the endogenous chelate in our
experiment may not be the same.
Prom the results of our own experiments as well as
those of others, we propose a modified mechanism of iron
transport and regulation, Figure 17. The flux of dietary
iron through the mucosal cells and into the blood is regu
lated by a series of equilibrium binding reactions. The
ligands of iron under normal conditions, are believed to be
compounds such as citrate, glucose and other common dietary
constituents, as well as compounds secreted by the mucosal
cells. The number of available specific sites on both
small and large molecules, as well as their individual
affinities for the iron manifest primary control. Further,
it is necessary that the iron be presented to the mucosal
cell as a soluble, preferably uncharged complex. Ingested
ionic iron, as it enters the lumen of the intestine and
the pH of its environment is raised, either combines with
available ligands provided by the diet or the seoretions
within the intestine, or else precipitates as an Insoluble
BLOOD EPITHELIAL CELL LUMEN
Other Low Molecular
Weight Ligands
Depot Cells
Ligands
Fe-Chelate— Transferrin Fe-Chelate ^ Pe-Chelate
Reticulocytes
Excretion
Macromolecules
Figure 17*— Proposed scheme for the mechanism of iron absorption from
the intestine by chelation.
CO
oo
89
hydroxide, carbonate, phytate, et cetera. Subsequent flux
of the Iron across cell nembranes, once it has entered the
circulation, involves participation of the complex metal.
It is important to realize that many low molecular weight
chelates may be operative in binding and transporting iron
in biological systems.
We have presented evidence that sugars under
specific conditions of concentration and pH, form stable
and soluble complexes with metal ions which would otherwise
be insoluble or sparingly soluble in an alkaline medium.
Large molar excess of sugar to metal ion favor chelate
formation. This finding explains the results of Bourne,
et al. (139) who studied complex formation with excess
metal ions added to small concentrations of polyol and
sugars. It is interesting to note that the complexes they
investigated contained molar ratios of metal to chelating
agent from 0.35 to 10.75. Their inability to show complex
formation of iron with glucose, and of cobalt and nickel
with dulcltol or glucose, can be attributed to both the
ratios of metal to sugar present and to their technique
of adding the metal salts to a pH 12 solution of the che
lating agent where the competition of hydroxide ion will
prevent the successful sequestering of iron by most ligands
for this metal.
Although chelation and complex formation of trace
90
metals with alcohols and polyols is often alluded to in the
literature, few, if any, definitive binding constants are
to be found. However we have been able to measure the
first binding constant of ferric fructose at acid pH.
Under the pH condition of the present study, the
reaction
H+ + (FeFr)2+^=± Fe3+ + Fr
is dominant, i.e., only one hydrogen ion is liberated per
mole of complex. At higher pHfs, up to three hydrogen ions
may be displaced. From steric considerations it is unlikely
that the fructose-iron can be more than bidentate. One of
these three protons probably comes from the acid dissocia
tion of co-ordinated water, as in the hydrolysis of the
free Fe3+. It therefore seems reasonable to postulate the
following sequences
[Fe(H20)i+(H2Fr)]3 + tFeCHgO)^(HFr) ]2+ + H+
[Fe(H20)i+(HFr)]2+ J&=± [Fe(H20)3(0H)(HFr) ]+ + H+
[Fe(H20)3(0H)(HFr)]+ [Fe(H20)3(0H)(Fr)] + H+
where fructose has been written as H2Fr to indicate that
two protons become acidic on co-ordination. The neutral
complex is quite stable, even in alkaline solutions,
although it does yield its metal rapidly to specific iron
91
binding sites of transferrin. This suggests that other
ohelating agents, of known equilibrium constants for com
plex formation with F*(jil)’ might be used as competitors
for fructose, thus allowing a direct determination of the
stability of the neutral fructose complex. We can estimate
-5
the product of to be about 10 , however, from the
requirement that the neutral form predominates at pH 3.5.
It Is interesting that the complexes of ?e(xu) ^
Ti(IV) glycerol and mannitol have been reported as
having a 1:1 ratio of metal to chelate. They are formed
with the dissociation of the proton from the polyol hydroxyl
group if the pH is in the range 1 to (lM)). Further
more, the reported absorption spectrum of the lronvmannltol
complex is very similar to that which we observed for iron-
sugar complexes, and the apparent dissociation constant for
the former was given as 0.35 at pH 2.5, In close agreement
with our value of 0.*+7 for the fructose-iron complex.
Various workers have studied the behavior of Fe^+
ions in water solution. Hedstrom (l*fl) found, by electro-
metric titration, that when the pH is Increased from low
values the hexaaquo ion at pH 1-2 starts transforming Into
a dimer
H *++
(H20 3^*. ^*e(H2° \
H
92
2+
In addition to the hydrolyzed species [(H20)jPeOH] and
[(HgO^PeCOH^]^*. The dimer has been shown by magnetic
susceptibility measurements, to be feebly paramagnetic or
diamagnetic (1^2). Proton relaxation studies (1M-3) were
consistent with this idea.
Martell, at *1. (lMf) proposed a dimeric structure
of copper-ATP chelate where two copper atoms are attached
through oxygen bridges. In general, it must be pointed out
that the presence of polymeric complexes are to be expected
but their ohemistry is quite complicated and poorly under*
stood (l*+5, l*+6).
Prom the results of our own experiments and those
of others we propose hypothetical structures for ferric
fructose as shown in Pigure 18. Figure l8a shows a possible
monomeric structure of ferric-fructose. However ESR and
NMR measurements of ferric fructose (R. Aasa, et al..
Biochim. Biophys. Acta in press) shows no spin signal and
the metal appears to have little influence on the alkyl
hydrogens from the fructose. Absence of spin signal is
indicative of iron-iron interaction which is only possible
if two iron atoms are within 7 1 of each other. This
leads us to postulate a dimeric configuration for ferric
fructose in which two iron atoms are linked to each other
by two ojQrgen bridges. Two possible dimeric configurations
for ferric fructose are shown in 18b and 18c.
93
HO
Figure 18.— Proposed structures of ferric fructose
chelate. Fructose in the structure is denoted by S.
9*f
Molecular weight determined by the freezing point
depression measurements Indicated ferric fructose to be a
low molecular weight compound. However, the presence of
sodium bicarbonate as an Impurity in the sample of ferric
fructose can seriously affect this measurement. In our
experimental conditions sodium bicarbonate could be a major
single source of contamination. Ferric fructose on elemen
tal analysis showed * + per cent sodium content. From the
titrimetric evidence of three hydrogen ions being liberated
for every Fe+++, the complex appears to have no charge.
Ferric fructose seemingly becomes negatively charged com
plex on electrophoresis (1^7) in bicarbonate buffer at pH
7.0. This would be expected if the complex travels as
bicarbonate bound complex.
/
A polymeric structure of ferric fructose shown in
Figure l8d appears to be consistent with the molecular
weight determined by ultracentrifugation experiments. Pre
liminary experiments with ultracentrifugation indicated a
small band of low molecular weight compounds moving along
with a large band of high molecular weight compounds. An
equilibrium existing between a polymer form and its smaller
units is also suggested by the dialysis experiment.
It appears likely that the ring configuration of
fructose participates in the chelation. Iron can have easy
access to both 1, 2- and 1, 6-position of fructose as can
95
be seen using Cortauldt atomic models. Several other posi
tions are available to consummate complexation with iron
but they are sterically unfavorable or hindered. Tosyla-
tion of the complex readily forms the 6-tosyl derivative
which suggests that the 6-position is not involved in che
lation. Thus we can predict a possible binding between
fructose and iron through 1, 2-position of fructose.
There are indications of different species of the
complex being formed with small changes in pH, molar con
centration and the rate of addition of alkali during the
preparation of these chelates. !Rrom the ultracentrifuga
tion experiments the presence of polymeric form of these
complexes have been indicated. However, these chelates
seem to be depolymerized in presence of excess ligands as
has been evidenced by the dialysis experiment. This physi
cal phenomenon can explain the fast uptake of iron in form
of ferric fructose when presented to the mucosal wall even
though it may be in polymeric form when administered.
Although sugar complexes have been reported pre
viously, very little has been known about their chemical
and physical properties. The work presented in this disser
tation Is only an initial approach to the understanding of
the chemistry as well as to the biological application of
metal sugar chelates.
CHAPTER V
SUMMARY
Evidence has been presented that the rate of iron
uptake both in vivo and in vitro is a function of the
chemical nature of the chelate used to complex the iron as
well as the concentration of complex in the intestinal
lumen. There is no indication of an active pumping mechan
ism or facilitated diffusion. The binding capacity of
transferrin does not affect the rate of entrance of iron
into the blood, although it does appear to control the
amount of iron which remains in the blood. The availabil
ity of the iron to tissues after absorption is also depend
ent upon the chelate used. Both ethylene-diaminetetra-
acetic acid and nitrilotriacetic acid complexes of iron
(Fe-EDTA and Fe-NTA) rapidly enter the blood but the
Fe-EDTA is immediately excreted through the kidneys whereas
Fe-NTA is deposited in the liver or bound to the trans
ferrin •
The molecular size and charge of complexes, as well
as their stability constants markedly affect the rate of
flux.
96
97
An endogenous chelate of low molecular weight has
been isolated from the intestinal lumen of the rabbit which
is able to complex iron. Participation of this endogenous
ligand has been suggested in the transport of iron through
the mucosal wall.
Sugars have been shown to form stable and soluble
complexes with metal ions under specific conditions of
concentration and pH. Evidence for the existence of the
complex is its solubility at alkaline pH, its characteris
tic spectral properties, titration and redox measurements
and direct isolation of the water soluble complex.
Ferric ion in the presence of fructose at pH 2.5
forms a complex which shows a distinct color change that
can be measured spectrophotometrically. The complex has
an absorption maximum at 370 m,M • From the variation of
the apparent extinction coefficient at 370 myu, with fruc
tose concentration, it has been determined that the complex
contains one fructose per iron and that the dissociation
constant is 0.^7 at pH 2.5.
Elemental analysis indicates the complex at pH 9.0
contains 2 Fe:2 fructose:1 Ha. The complex may exist in
polymeric forms as determined by ultracentrifugation, how
ever, the dimer is in rapid equilibrium with the polymer in
presence of excess sugar. Determination of the structure
of ferric fructose has utilized hydrolytic studies of the
98
complex, esterification of the complex by tosylatlon, redox
and titrimetric measurements, ionophoresis and infra-red
studies. An attempt has been made to elucidate the struc
ture of metal sugar complexes by partial periodation tech
nique. Based on these experiments, possible dimeric and
polymeric structures of ferric fructose are proposed.
LITERATURE CITED
LITERATURE CITED
1. Hahn, P. P., Bale, W. P., Be If our, W. M., and
Whipple, A. H., J. Exper. Med,, 2£Li lo9 (19*+3)«
2. Granick, S., Bull. New York Acad. Med., 30* 81 (195*0.
3. Bale, W. P., Roes, J. P.. and Belfour, M. M.,
J. Exper. Med., £8, 169 (19*+3).
if. Granick, S., J. Biol. Chem., i6*f, 737 (19W.
5. Gabrio, B. W., and Salomon. K., Proc. Soc. Exper.
Biol. Med., 2ly 12*fr (195b).
6. Brown, E. B., and Bother, M. L., J. Lab. & Clin. Med.,
6g, 357 (1963).
7. Bothwell, T. H., Pirozio-Biroli and Pinch, C. A.,
J. Lab. & Clin. Med., Jjl, 2*f (1958).
8. Dowdle, E. B., Schachter, D., and Schenker, H.,
Am. J. Physiol., 198. 609 (i960).
9. Moore, C. V., The Harvey Lectures, 67 (1961).
10. Charley, P. J., Stitt, C., Shore, E., and Saltman, P.,
J. Lab. & Clin. Med., 61, 397 (19b3).
11. Josephs, H. W., Blood, 13, 1 (1958).
12. Moore, C. V., Dubach, R,, Minnich, V., and
Roberts, H. K., J. Clin. Invest., 23., 755 (19***0 •
13. Greenberg, S. M., Tucker, R. G., Heming, A. E., and
Mathues, J. K., J. Nutrition, 63, 19 (1957).
l*f. Gubler, C. J., Science, 123. 87 (1956).
15. Groen, J., Van Den Broek, W. A., and Veldman, H.,
Biochim. et Biophys. Acta, 1, 315 (19^7).
16. Woodruff, C. E., Borden’s Rev. Nutritional Research,
20, 61 (1959;.
100
101
17. Kovacs. E., and Bemat) J,, Therapia Hungarica, j>,
1 (1957).
18* franklin, M., Rohse, W. G., de la Huerga. J.. and
Kemp, C. R., J. Amer. Med. Aasoe., l6o. 1d85 (1958X
19. Herndon, J. P., Rice, E. 0., Tucker, R. G.,
Van Loon, E. M«, and Greenberg, S. M., J. Nutri
tion, 6}f, 615 (1958).
20. Conrad, M. E., and Crosby, W. H., Blood, 22, M-06
(1963).
21. field, M., Sekl, M., Mitchell, M. L., and
Chambers, T. C., J. Lab. < S c Clin. Med., 5£, 929
(I960).
22. Pollack, S., Kaufman, R. M., and Crosby, W. H.,
Science, lMt, 1015 (196*f).
23. Wheby, M. S., Conrad, J. E., Heldberg, S. E.. and
Crosby, W. H., Gastroenterology, *f2, 319 (1962).
2b, Mendel, G. A., J. Lab. & Clin. Med., ^8, 9^3 (1961).
25. Wheby, M. S., and Crosby, W. H., Blood, 22, *fl6
(1963).
26. Perrin, D. D. , J. Chem. Soc., Part III, 3120 (1958).
27. Perrin, D. D., J. Chem. Soc., Part III, 3125 (1958).
28. Clancy, J. B., Aldrich, R. H. Rummel, W., and
Candon, B. H., Am. Pract. & Digest Treat., 8,
19^8 (1957).
29. Beutler, E., and harsh, S. E., New England J. Med.,
267. 538 (1962).
30. Jacobi, H., Pfleger, K., and Rummel, W., Arch. Exper.
Path, u Pharmakol., 229. 198 (1956).
31. Saltman, P., fiskin, R. D., and Bellinger, S. B.,
J. Biol. Chem., 220, 7bl (1956).
32. Saltman, P., fiskin. R. D., Bellinger, S. B., and
Alex, T., J. Biol. Chem., 220. 751 (1956).
33* Saltman, P., frisch, H., fiskin, R. D., and Alex, T.,
J. Biol. Chem., M , 71 fl (19&).
102
3*t* Bass, R., Bemick, S.. and Saltman, P., Exper. Cell
Research, 1^, 395 (1959).
35* Brown, E. B., Jr., and Justus, B. W., Am. J. Physiol.,
I2i> 319 (1958).
36. Cartwright, E. G., Wintrobe, M. M., J. Clin. Invest.,
2£, 86 (19^9)•
37. Rath, C. E., and Pinch, C. A., J. Clin. Invest., 28.
79 (19^9).
38. Ehrenberg, A., and Laurell, C. B., Acta. Chem. Scand.,
2, 68 (1955).
39. Wallenius, G., Scand. J. Clin, and Lab. Invest., *f,
2»f (1952).
1+0■ Giblet, E. R., Hickman, C. G., and Smithies, 0.,
Nature, l8l» !589 (1959).
1+1. Dera, R. J., Alberto, M., and Glynn, M. P., J. Lab. &
Clin. Med., 6l, 280 (1963).
1+2. Pollack, S., Balcerzak, S. P., and Crosby, W. H.,
Blood, 21, 33 (1963).
1+3. Hallberg, L., and Solvell, L., Acta Med. Scandinav,
Suppl., 358 (I960).
M+. Pirzio-Biroli, J. Lab. & Clin. Med., _gL, 37 (1958).
1+5. Heilmeyer, L., Keller, W., Vivell, 0., Kiederling, W.,
BeIke, K., Wohler, P., and Schultz. H. E.,
Deutsches Med. Wochnschr., 6, 385 (1961).
1+6. Aasa, R., Malmstrom, B., Saltman, P., and Vanngard, T.,
Biochim. et Biophys. Acta, 25* 203 (1963).
1+7. Shade, A., Reinhart, R., and Levy, H., Arch. Biochem.,
20, 170 (19^9).
1+8. Laurell, C. B., Pharmacol. Rev., J+, 371 (1952).
1+9. Jandle, J. H., Imman, J. K., Simmons, R. L.. and
Allen, D. W., J. Clin. Invest., ^o, l6l (1959).
50. Michaelis, L., Adv. in Protein Chemistry, 2* 53 (19^71
51. Bielig, H., Bayer, E., Naturwiss., ^f2, 125 (1955).
103
52. Loevus, M., and Fineberg, R., Biochlm. et Biophys.
Acta, 26, Mfl (1952).
53. Mazur, A., and Carleton, A., J. Biol. Chem., 236,
1817 (1963).
5*f. Mazur, A., Baez, S., and Shorr, E., J. Biol. Chem.,
213. lb? (1955).
55. Mazur, A., in "Metaholism and Function of Iron,"
Report of Nineteenth Ross Pediatric Research Con
ference, p. 27, Columbus, Ohio (1956).
56. Green, S., and Mazur, A. J. Biol. Chem., 227. 653
(1957).
57. Morgan, G. T., and Brew, H. D. K., J. Chem. Soc.,
121. l^f56 (1922).
56. Charley, P., Rosenatein, M., Shore, E., and
Saltman, P., Arch. Biochem. and Biophys., 88,
222 (I960).
59. Fisher, H., and Orth, H., "Bie Chemie des Pyrroles."
II Band I Halfte, Leipzig, Akademische
Verlagsgesellschft (1937).
60. Michaelis, S., and Hahn, P. F., J. Biol. Chem., 155.
661 (19M+).
61. Pirzio-Biroli, G*, in "Metabolism and Function of
Iron," Report of the Nineteenth Ross Pediatric
Research Conference, p. 2*+, Columbus, Ohio (1956).
62. Wacker, W. E. C., and Vallee, B. L., J. Biol. Chem.,
23*f. 3257 (1959).
63. Martell, A. E., and Calvin, "Chemistry of the Metal
Chelate Compounds," p. 399, Prentice-Hall, Inc.,
Englewood Cliffs, N. J. (1952).
6*+. Warner, R., and Weber, I., J. Am. Chem. Soc., 75.
5086 (1953).
65. Hamm, R. E., Shull, C. M., Jr., and Grant, B. M.,
J. Am. Chem. Soc., 2111 (1952).
66. Webster, 0, W., Mahler, W., and Benson, R. E.,
J. Am. Chem. Soc., 8Jf, 3678 (1962).
10*f
67. Corbett, J. D., Druding, L. F., Burkhard, W. J., and
Lindahl, C. B., Discussions Farady Soc., 32, 79
(1961).
68. Corbett, J. D., Inorg. Chem., 1, 700 (1962).
69. Ray, P.. and Sen, A.,J. Indian Chem. Soc., 25, *+73
(19W.
70. Anderegg, G., Flaschlka, H., Sallmann, T„, and
Schwarzenbach, G., Helv. Chim. Acta, 32, 113 (1951 *).
71. Schwarzenbach, G., and Magyr, B., Helv. Chim. Acta,
}f5, 11*25 (1962).
72. Forth, W., and Seifen, E., Arch. Exptl. Pathol.
Pharmakol., 2M-1. 556 (1961).
73* Brise, H., and Hallberg, L., Acta Med. Scand., 171.
Suppl. 376, 51 (1962;
7**. Rummel, W., and Candon, B., Intern. Record Med. &
Gen. Praot. Clin., 162, 783 (1956).
75. Hopping, J. M., and Rullifsen, W. S., Am. J. Phys.,
205, 885 (1963).
76. Nissim, J. A., J. Physiol., 118 (Suppl. b), 6b (1962).
77. Lange, W. E., Barnet, A. G., and Foye, W. 0,,
J. Pharm. Sci., j?l, 1128 (1962).
78. Rubin, M., and Princiotto, J. V., J. Agr. Food Chem.,
11, 98 (1963).
79. Nielands, J. B., Bact. Rev., 21, 101 (1957).
80. Tiffin, L. 0., and Brown, J. C., Science, 135. 311
( 1962) .
81. Charberk, S., and Martell, A. E., "Organic Sequester
ing Agents," p. *+l6, John Wiley & Sons, Inc.,
New York (1959).
82. Leonard, C. D., and Stewart, I., Proc. Florida State
Hort. Soc., 6*f, 20 (1952;.
83. Leonard, C. D., and Stewart, I., Proc. Florida State
Hort. Soc., 66, **9 (1951 *;.
8*t.
85.
86.
87.
88.
89.
90.
91.
92.
93.
9**.
95.
96.
97.
98.
99.
100.
101.
102.
105
Stewart. I., and Leonard, C. D., Science, 116.
56*+ (1952).
Reitz, H. J., and Leonard, C. D., Florida Univ. Agr.
Expt. Sta. (Gainesville) Bull. 536 (195*0.
Wallace, A., North, C. P., Kofranek, and Lunt, 0. R.,
Calif. Agr., 2, 13 (1953).
Hilgeman, R., Progr. Agr. Ariz., 8, 1 (1956).
Wallace, A., Ed., "Symposium on the Use of Metal
Chelates in Plant Nutrition," The National Press,
Palo Alto, Calif. (1956).
Gerhard, P. B., Am. Rose Mag., 2£ (July 1955).
Beattie, J. M., Proc. Ohio State Hort. Soc., 175
(195*0.
Waters, L. L., and Stock, C., Science, 102. 601
(19*0).
Kety, S. S., J. Biol. Chem., 1^+2, 181 (19*+2).
Schubert, J., Science, 105. 389 (19*+7).
Schubert, J., Ann. Rev. Nuclear Science, j>, 369 (1955)
Foreman, H., Huff, R. L., Oda, J. M., and Garcia, J.,
Proc. Soc. Exp. Biol, and Med., 22, 520 (1952).
Tripod, J., and Keberle, A., Helv. Phys. Pharmacol.
Acta, 20, 291 (1962).
Heilmeyer, L., and Wohler, P., Deut. Med. Wochechr.,
82, 2661 (1962).
Spencer, H., Feldstein, A., and Smachson, J.,
J. Lab. & Clin. Med., M+5 (1962).
Larsen, B. A., Bidwell, R. G. S., and Howkins, W. W.,
Can. J. Biochem. ana Physiol., 38, 51 (I960).
Mills, C. F*, Biochem. J., 63. , 187 (1956).
Mills, C. F., Biochem. J., 6£, 190 (1956).
Wallerstein, R. 0., and Mettier, S. M., eds., "Iron
in Clinical Medicine," University or Calif. Press,
Berkeley (19 58).
103.
10k.
105.
106.
107.
108.
109.
110.
111.
112.
113.
ll*t.
115.
116.
117.
118.
119.
120.
106
Cotzias, G. C., Physiol. Revs., ^8, 503 (1958).
Peters, T., Giovannlello, T. J., Apt, L., and
Ross, J. P., J. Lab. A Chem. Med., V8, 280 (1956).
Traube, W., Kuhbier, P., and Harting, H., Ber., 66,
19^5 (1933).
Leiser, T., Ebert, R., Ann. N. Y. Acad. Sci. U. S.,
532. 89 (1937).
Bourne, E. J., Nery, R., and Weiger, H., Chemistry
and Industry, 998 (1959).
Slack, C. G. B., and Wilkinson, J. P., Lancet, 1, 11
(1W).
Nissin, J. A., and Robson, J. M. , Lancet, 1, 686
(19^9).
Bersin, T., Pharm. Acta Helv., 26, *f07 (1951).
Peligot, E., Compt. rend., 20s 153 (1880).
Winter, H., Ann., 2Mf. 295 (1888).
Jackson, R. P., Silsbee, C. G., and Proffit, M. J.,
Bureau of Standards Sci. Papers, 20, 586 (1926).
Mills, J. A., Biochem. Biophys. Research Commons.,
6, ^18 ( 1961) .
Charley, P., and Saltman, P., Science, 139. 1205
(1963).
Roe, J. H,, J. Biol. Chem., 102, 15 (193*0.
Nakamoto, K.. "Infrared Spectra of Inorganic and
Coordination Compounds," p. 156, John Wiley and
Sons, Inc., New York (1963).
Pujita, J., Nakamoto, K., and Kobayashi, M.,
J. Am. Chem. Soc., 20s 3963 (1956).
Umbreit, W. W., Burris, R. H., and Stauffer, J. P.,
"Manometric Techniques," Burgess Publishing Co.,
Minneapolis (1957).
Lieser, T., and Schweizer, R., J. Am. Chem. Soc.,
519. 271 (1935).
121. Compton, J., J. Am. Chem. Soc., 60, 395 (1938)*
122. James, B. R., Iyons, J. R., and Williams, R. J. P.,
Biochemistry, 1, 379 (1982).
123. Hi 1 bum, R, M., and Vosburgh, W. C., J. Am. Chem.
Soc., 2Z» 1352 (1955).
12^. Malaprade, L., Bull. Soc. Chim. Prance, 883
( 1928) .
125. Malaprade, L., Compt. rend, 186. 382 (1928).
126. Schwarz, J. C. P., J. Chem. Soc., 276 (1957).
127. Barker, S. A., and Bourne, E. J., J. Chem. Soc.,
905 (1952).
128. Courtois, A., and Guemet, B. G., Bull. Soc. Chim.
Prance, 1388 (1957).
129. Zuman, T., Sicher, P., Kupricka, B.. and Svoboda, H.,
Coll. Czech. Chem. Comm., 1388 (1958).
130. Bobbit, J., Adv. Carbohydrate Chem., 11, 1 (1956).
131. Keen, J., Symons, P., Proc. Chem. Soc., 383 (I960).
132. Buist, G. J., Bunton, C. A., and Miles, J. H.,
J. Chem. Soc., b567 (1957), 7^3 (1959).
133. Hutson, D. H., and Weigel, H., J. Chem. Soc., 15^6
(1961).
13^. Poster, A. B., Adv. Carbohydrate Chem., 12, 81
(1957). “
135. B©lli N., Palmer, T., and Johns, P., J. Chem. Soc.,
1536 (19^9).
136. Bell, N., Greville, S., J. Chem. Soc., 1902 (1950).
137. Dawson, R. M. C., Elliott, D. C., Elliott, W. H.,
Jones, K. M., "Data for Biochemical Research,"
p. 23o, Oxford University Press, London (1959).
138. Eppley, R. W., Gee, R., and Saltman, P., Physiol.
Plant., 16, 777 (1963).
139. Bourne, E. J., Nery, R., and Weiger, H., Chemistry
and Industry, 998 (1959).
I»f0.
l*fl.
Ilf 2.
Ilf 3.
lMf.
I*f5.
li+6.
1»+7.
108
Pyatnitsky, I. V*. and KZLibua, A. K., Ukrainian Chem*
Jour., 22, MtO (1963).
Hedstrom, B. 0. A., Arkiv. Kami, 6, 1 (1953).
Malay, I». N., and Naylor, M. C., In S. Kirschner,
"Advances in the Chemistry of the Coordination
Compounds," p. 520, Macmillan Co., New Tork (1961).
Broersma, S., J. Chem. Phys., 26, l*+05 (1957).
Taqui Khan. M. M., and Martell, A. E. , J. Phys. Chem.,
66, 10 ( 1962) .
Tobias, R. S., and Tyree, S. Y., J. Am. Chem. Soc.,
§1/6385 (1959).
Tobias, R. S., and Tyree, S. Y., J. Am. Chem. Soc.,
82, 32¥+ (I960).
Charley, P., Sarkar, B., Stitt, C., and Saltman, P.,
Biochim. Biophys. Acta, 6£, 313 (1963).
T h is d is s e r ta tio n h a s b e e n 65— 3116
m ic r o film e d e x a c tly a s r e c e iv e d
SA RK AR, B ib u d h en d ra, 1 9 3 5 -
ST U D IE S O N THE T R A N SP O R T , M ETABO LISM
AND CH EM ISTRY O F IR O N -SU G A R C H E L A T E S.
U n iv e r s ity o f S ou th ern C a lifo r n ia , P h .D ., 1964
C h e m is tr y , b io lo g ic a l
University Microfilms, Inc., Ann Arbor, Michigan

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Asset Metadata

Creator Sarkar, Bibudhendra, 1935- (author)

Core Title Studies on the transport, metabolism and chemistry of iron-sugar chelates

School Graduate School

Degree Doctor of Philosophy

Degree Program Biochemistry

Degree Conferral Date 1964-06

Publisher University of Southern California (original), University of Southern California. Libraries (digital)

Tag chemistry, biochemistry,OAI-PMH Harvest

Language English

Contributor Digitized by ProQuest (provenance)

Advisor Saltman, Paul (committee chair), Adamson, Arthur W. (committee member), Fluharty, Arvin L. (committee member), Mehl, John W. (committee member)

Permanent Link (DOI) https://doi.org/10.25549/usctheses-c18-362216

Unique identifier UC11359213

Identifier 6503116.pdf (filename),usctheses-c18-362216 (legacy record id)

Legacy Identifier 6503116.pdf

Dmrecord 362216

Document Type Dissertation

Rights Sarkar, Bibudhendra

Type texts

Source University of Southern California (contributing entity), University of Southern California Dissertations and Theses (collection)

Access Conditions The author retains rights to his/her dissertation, thesis or other graduate work according to U.S. copyright law. Electronic access is being provided by the USC Libraries in agreement with the au...

Repository Name University of Southern California Digital Library

Repository Location USC Digital Library, University of Southern California, University Park Campus, Los Angeles, California 90089, USA