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~sliue D a 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
