A Laboratory Experiment Illustrating the Properties
and Bioavailability of lron
Doris R. ~imbrou~h,'
Noelia ~artinez?and Stephanie stolfusz
Department of Chemistry, Box 194,P.O. Box 173364,University of Colorado at Denver, Denver, CO 80217-3364
Many students that take chemistry do so because chemistry is required to pursue assorted careers in the health
sciences: nursing, medicine, dentistry, pharmacy, etc.
These students a r e often less t h a n enthusiastic about
learning in the nonbiologically oriented subdisciplines of
chemistry, such a s inorganic, analytical, and physical.
Adding a hiological "spin" can make these topics more appealing to those students who do not find them intrinsically interesting.
Included i n this paper is a description of a qualitative
laboratory experiment designed for either general chemistry or prenursinglprephysical therapy chemistry courses.
It could also be easily adapted for a high school chemistry
project. I t relates the properties and behavior of iron and
its complexes to the absorption and bioavailablity of iron in
human systems.
lron Absorption and Function in the Body
Iron is the most abundant trace element in the human
body (I1. I t is found circulating in the blood in hemoglobin
and transferrin, and in muscle tissue in myoglobin. I t is
stored in the liver, spleen, and bone marrow in ferritin and
is associated with various enzymes (1,2).Most of us absorb
less than 15% of the 10-15 mg of iron contained i n our
daily diets, in part because not all of the iron that we ingest
is bioavailable; our bodies are unable to absorb iron in
some forms.
Iron absorption i s also limited by a control mechanism in
our intestines, where absorption occurs. This control
mechanism, which is not well-understood, decreases the
uptake of iron when the body's needs have been met ( 1 , 3 ) .
When the body's needs have not been met, i n a condition
called anemia, a normally functioning body will respond by
increasing its uptake of iron to 20 or 30% of what is ingested (I).Increased absorption occurs commonly during
pregnancy or when a large number of red blood cells have
been lost through bleeding. Conversely, a n excess of iron i n
our systems, a much more rare condition called hemochromatosis, is a toxic condition that i s detrimental to liver,
heart, and pancreas function (1).
Many foods are enriched by the manufacturer with iron
as a nutrient. Most consumers, even chemists, are often
surprised to learn that food labels listing "iron" or "reduced
irou" mean t h e foods have actual metallic iron filings
added to them. The hioavailability of iron in this form is
open to question, and pH a s well a s the complexes that irou
fonns with other nutrients affect whether this iron is ahsorbed.
Iron in the metallic form cannot be absorbed, but presumably it dissolves a s i t is oxidized by the hydrochloric
acid in the pH 2 environment of the stomach. However,
iron absorption occurs i n t h e small intestine, which is
strongly buffered to pH 8. Both Fe(I1) and Fe(II1) precipitate a s hydroxides a t basic pH, and in order to be absorbed
'Author to whom all correspondence should be addressed.
2High school student participating in a mentorship program for female high school students.
558
Journal of Chemical Education
Substances That Enhance and Inhibit the Absorption of
lron in the Small Intestine ( 1 )
Substance
Enhance Uptake
ascorbic acid (vitamin C)
citric acid
fructose
histidine
lysine
methionine
Result after raising pH > 8
clear, pale green solution
clear, almost colorless solution
clear, yellow-green solution
green solution with trace of
brown precipitate
clear, yellow solution
orange solution with trace
of orange precipitate
inhibitors
carbonate (NazC03)
greenish black precipitate
phosphate (NasPOa)
greenish black precipitate
oxalic acid
deep brown solution
Each substance was added in excess to a solution of iron at pH 2.
and then the pH was increased to 8 or above by the addition of a 5%
solution of NaOH. In the case of the sodium carbonate and sodium
phosphate, additional base was not required to raise the pH.
the iron must be soluble. Many substances complex with or
chelate iron and k e e i~t soluble. assisting with its absomtion (I). Some of these substances are lGted i n the tabie.
Conversely, there are several substances commonly present i n food digestion that hinder iron absorption presumably because they enhance iron precipitation ( I ) .
Experimental
Caution C a r t must bc used in hond~nghydn,rhloric and
w 4 l r art& and in ywpnrinp t h e 5 ' ~adcu~nn).dn,xdt 4 u -
Isolation of Metallic lron from Cereal
Iron was isolated from Total cereal using a modified version of a previously described method (4,5). From approximately 28 g of cereal (1serving), a n easily visible amount
of iron filings can be isolated. Grinding the cereal with the
bottom of a n Erlenmeyer flask on top of plain white paper,
a small amount a t a time, was found to be more efficient
than using a mortar and pestle. The dry, ground-up cereal
is transferred to a 250-mL beaker and stirred with a cow
magnet (4). Small amounts of cereal stick to the magnet
and are brushed off into another smaller beaker (100 or 50
mL). This process is repeated until significantly less cereal
sticks to the magnet when the crushed cereal is stirred;
most of the iron is isolated within 20 min.
Enough water i s added to the second beaker to cover the
cereal crumbs separated with the magnet by approximately 1cm. The magnet is then held against the bottom
of the outside of the beaker a s the contents of the beaker
are swirled. Within a few seconds of swirling. iron filings
are visible near the magnet. These can be isolated by using
the maenet on the outside of the beaker to gentlv
. move the
filings tb the side of the beaker and up, away from the wet
cereal. There they can be captured by a spatula.
Behavior of Iron in Simulated Physiological Conditions
Commercial iron filings are used in this portion, a s the
iron isolated in the above procedure is not sufficient to prepare a solution concentrated enough to clearly observe the
effects described.
Stomach
Iron metal can he dissolved i n a pH 2 hvdrochloric
arid potassium chlor~debufrer suluriun that ;i;nulates the
m v i r o n m c ~of
~ tthe stomach. Tht: inm dis;olves as it is oxidized to iron(I1) and hydrogen gas evolves. Approximately
50-100 mg of iron filings can be dissolved in 125 mL of the
buffer solution (250 mL 0.20 M KC1 solution and 65 mL
0.20 M HC1 solution, diluted to 1.00 L). Although one can
observe the hydrogen gas evolution almost immediately,
the dissolution requires several hours and is best accomplished overnight or over a week's time.
Intestine
A small portion of the pH 2 solution of iron prepared
above (5-10 mL) is placed into a test tube. A5% s&iion of
sodium hydroxide is added dropwise to raise the pH to approximately 8, thus sirnulatinithe pnisngc from thc acidic
iromuch to the strongly buffered haiic small intestme. The
o H is -monitored
ci~loruHast(EMScience', indicator
~ ~ ~ - uiine
- ~
strips. Regular pH paper was somewhat inaccurate; a pH
meter can also be used but is cumbersome. As the DHrises.
the solution turns greenish brown and a dark-brown, almost black...precipitate
. forms. This precipitate
. . can he isolated and will not completely redissolve upon acidification
of the solution.
Thc precip~tatecan be nvoided by the addition of jeverr~l
rhelatjne anent.; which arc listed in the tahlr. I k t rc;iults
are obtained when the chelating agent is in large molar
excess of the iron. To 5-10 mL of the pH 2 solution of the
iron prepared above, approximately 0.5 g (a scoopula) of
the chelating agent is added. The test tube contents are
stirred or shaken to dissolve the chelating agent. The pH
is then increased to 8 by the dropwise addition of the 5%
sodium hydroxide solution. If enough chelating agent has
been added, the solution will remain clear of precipitate,
although it does often turn yellow or green as the pH is
increased. Solutions of oxalic acid-iron complex turn a
deep brown.
L ~ - - - ~
~
~
-
- -
Discussion
Presumahlv metallic iron dissolves in the stomach as it
reacts with k e hydrochloric acid. This is simulated in the
above process a s iron is dissolved in the solution of pH 2.
Fds)+ ZHCl(aq)4 ~ e "+ 2C1-+ Hz
I n the absence of chelating agents, increasing the pH
causes the oxidation of iron(I1) to iron(II1) and the formation of colloidal gels with the ultimate precipitation of the
iron(II1) hydroxide (6).As the pH is increased, precipitation is often preceded by a yellow color that is caused by
the formation of hydrolyzed solvated ion species (6).
The formation of the precipitate dramaticallv illustrates
the biounavailahility df iron in the absence bf chelating
agents. The precipitate
would he excreted through the
.
.
large intestine, as is the case with excess dietary or supplementaw iron that is not needed by the body.
As the table shows, lntestmal iron absorption is hindercd bv the Dresence o l ' ~ h o s ~ h aund
t r carbonat~s.Addition of sodium carbonateAortiibasic sodium phosphate to
the pH 2 iron solution produces a n increase in pH analogous to the addition of 5% sodium hydroxide, with the subsequent formation of precipitate indicating biounavailahility.
Addition of a chelatine aeent com~lexesthe iron so that
it can withstand the increase in pH and remain in solution.
The better the iron is chelated, the better are the chances
for its absorption. We found that fructose and citric acid
acted as the best chelating agents. Ascorbic acid was less
effective, and the amino acids were the least effective. All
of the substances worked to some extent, and the difference in behavior of the solutions with and without the
presence of the chelating agent as the pH was increased
was significant. This simulation effectively provides the
student with the understanding of the vital importance of
dietarv chelatine aeents for the bioabsor~tionof iron.
while "simultane&sG illustrating the properties of iron;
oxidation-reduction chemistnr, and the solubilitv differences that can be brought aboit by complexation.
As can be seen in the table, oxalic acid hinders the intestinal uptake of iron. Spinach contains a high concentration
of oxalates. and for this reason. much of the iron in soinach
is biounavailahle ( 1 ) .When oxdic acid is added to {he pH
2 solution of iron and the pH is subsequently increased,
the solution turns a deep brown, but no precipitate forms.
The iron-oxalate comolex is ~resumablvnot absorbed for
reasons other than the precipitation discussed above, and
this comolexitv is not demonstrated bv this s i m ~ l i s t i c
model of the digestive system.
~~~~
~~
Acknowledgment
We are grateful to Corime Campbell and Joel Selbin for
their help in understanding the biochemical uptake and
chemical properties of iron. Funding for this project was provided as part of the CU-Denver Female Mentorship in Science and Technology Program with a grant from the Colorado
Community College and Occupational Education System,
Federal Vocational Education Discretionary Fund.
Literature Cited
1. Linder M. C. InNutritionolBiocharnislryand
Metabolism vilh ClinicolAppiimtions;
Lindcr, M. C., Ed.;Elseuier: NewYork, 1985.
2. Jones, M. M.:Johnston.D. 0.: Nettelville, J.T.; Wmd. J. L.; Joesten, M. D.Ch~rn~stry
ond Society; Saunden: Philadelphia, 1987.
3. Davenport, H. W. Physiology oflhe D ~ g ~ s l i u
Dea d , 5th ed.; Year Book Medical: Chi-
6. Greenwmd,N. N.; Earn8haw.A. Chemistry oftheElernmfs: Pergamon: Oxford, 1984.
3Address: 480 Democrat Road, Gibbstown, NJ 08027, an associate of Merck.
Volume 72 Number 6 June 1995
559