"In honor of the International Year of the Periodic Table this series of articles details the Element of the Month project developed by Stephen W. Wright (SWW), Associate Research Fellow at Pfizer Inc., and Marsha R. Folger (MRF), chemistry teacher (now retired) at Lyme – Old Lyme High School in Connecticut. Read The Element of the Month - An Introduction for an overview of the project and links to the other articles in the series." - Editor
Iron is presented in May as the last Element of the Month in our program. The students are familiar with iron and most will know that steel is primarily composed of iron. All students are familiar with steel and the endless variety of items fabricated from it. And all students will be very aware of the tendency of iron and steel to oxidize to form rust. However, the class will generally not be aware of how much the oxidation of iron costs the global economy and they will mostly be unaware of the strategies used to prevent rusting of iron and steel objects. The chemistry of iron is very colorful, and iron forms a nearly endless array of colored complexes and precipitates. For the last Element of the Month, we discuss the chemistry of iron and perform demonstrations as were done with the previous elements, and we end with a hands-on lab activity that the entire class can participate in.
Occurrence in Nature
Iron is abundant in the earth’s crust. By mass it is the most common element on Earth. It occurs in innumerable minerals containing silicon, aluminum and oxygen, and frequently combined with other metals such as calcium and magnesium. The best known iron containing minerals are pyrite (fool’s gold, FeS2), and the ores limonite (FeO(OH) · n H2O), magnetite (Fe3O4), and hematite (Fe2O3). Magnetite is frequently found to be naturally magnetic. It was once known as lodestone and was used to make the first magnetic compasses for navigation. Iron oxides are usually responsible for the red brown color on weathered rocks. Even more iron is to be found in the earth’s mantle and core. Iron is the most common transition metal, ten times more common than all the other transition metals combined. Rarely, iron is found as the free element in nature in meteorite fragments. We pass around samples of magnetite, pyrite, and hematite for the students to examine.
Iron is essential for life. It is an important component of various proteins in the body, most notably the red oxygen carrying protein hemoglobin. Iron and its steel alloys are essential for modern societies. We show on the lecture desk a box of nails, a short length of railroad iron, a hammer, a wrench, and a box of an iron - fortified cereal. We ask the class how iron is prepared and generally the class will quickly answer that it is prepared by the reduction of iron oxides by carbon at very high temperatures.
Figure 1: Left-test tubes showing colors of FeCl3, KSCN and K4Fe(CN)6, Right-colors after adding FeCl3 to other solutions
Iron is a relatively dense metal by comparison to other metals that the students have encountered. It is malleable, meaning it can be shaped by hammering, but it is still a hard metal that makes it useful for the manufacture of tools. It has a relatively high melting point (1538 °C, or about 2800 °F). Iron conducts electricity but it is not an efficient conductor and therefore sees no use for this purpose.
Figure 2: Left - potassium thiocyanate & potassium ferrocyanide solutions painted used to paint IRON on paper, Right - writing is sprayed with iron III chloride
Iron is a moderately reactive metal, with three important oxidation states: 0, +2, and +3. We recall for the class the copper Element of the Month experiments and note that iron dissolves in acids, and is more easily oxidized than copper and silver but is less easily oxidized than zinc, magnesium, or aluminum. We place a 7 to 10 cm length of clean iron wire in a test tube, cover it with 6 M HCl, and briefly warm the test tube in beaker of hot water.1 Next we note that there are two diagnostic tests for iron, and we show the reaction of potassium thiocyanate (KSCN) solution and potassium ferrocyanide (K4Fe(CN)6) solution with a 5% (w/v) ferric chloride (FeCl3) solution (figure 1).2 Next, we perform a bit of chemical magic. To do this, we spray dilute FeCl3 solution on a sheet of paper on which a picture has been painted using potassium thiocyanate solution and potassium ferrocyanide solution and allowed to dry (figure 2).3 We note that iron forms many colored complexes, just as copper did, which is typical of the transition metals. We perform another chemical magic trick, the “Jug of Mystery” demonstration, and show the class some of the colors that can be produced from iron compounds (figure 3).4 As an essential nutrient, iron is added to foods to supply “100% of the recommended daily allowance” of iron. We add a cup or so of an iron – fortified breakfast cereal to water in a beaker, show a clean magnetic stir bar to the class, add the stir bar to the beaker, and place the mixture on a magnetic stirrer. We note that the ingredients panel on the cereal box states that the iron is present as iron itself and not as an iron compound, and ask the class "why is the iron is added to the cereal as metallic iron?" The answer lies in the chemistry of iron: humans can absorb iron (II) compounds but do not absorb iron (III) compounds well. However, if an iron (II) compound was added to the cereal, it would react with oxygen in the air to yield an iron (III) compound. Therefore, the iron is added as iron metal, which dissolves in the HCl in the stomach to afford iron (II) chloride, which is quickly absorbed.
Figure 3: Jug of Mystery demo4
The lowest cost and most frequently used method to retard the oxidation of objects made from iron is to coat the iron with a physical barrier to exclude air and moisture. Most frequently the physical barrier is either paint or a film of oil. While inexpensive, the disadvantage of this method is that the coating is not actually impermeable to oxygen and moisture, and the iron underneath will be exposed if the coating is damaged in any way. Hence frequent maintenance of the coating is required to preserve the metal underneath.
A more effective, and more expensive, method to slow down the oxidation of an iron object is to place it in electrical contact with a metal that is more reactive and has a higher oxidation potential than iron. Such a “sacrificial” metal will react with oxygen in preference to the iron which it is in contact with. We remind the class about the relative activities of metals that we observed during our experiments with copper. We ask the class "what metal might be used as the sacrificial metal?" Zinc, aluminum and magnesium are frequently offered as suggestions.5 Zinc is most frequently used because it is inexpensive and bonds well to iron. Items treated in this way are known as galvanized. Magnesium is used in some situations, for example to protect the hulls of deep-water ships. As a thought experiment, we ask "what would happen if iron was in contact with a metal less reactive than iron, for example copper?" In this case the iron becomes the sacrificial metal! This phenomenon is called galvanic corrosion. People who are familiar with watercraft and ships know to avoid such situations.
The third, and most expensive technique, is to alloy the iron with other metals, primarily chromium. This is how “stainless steel” is made. There is enough chromium in the metal to form a protective surface film of chromium oxide upon reaction with the oxygen in air. Like aluminum, this tough oxide layer adheres to the metal and prevents further oxidation by preventing oxygen from accessing the metal underneath. The resistance of a particular stainless steel alloy to corrosion depends upon the chemical composition of that stainless steel alloy.
We ask the class "what factors speed up the rusting of iron and steel?" Water, salt, and acids are usually correctly offered as answers. At this point, we bring the students to their lab benches and we guide them through a hands – on activity to study of the corrosion of iron by exposing steel nails to various conditions in petri dishes containing a gelatin matrix.6 Before the class departs, we check the cereal mixture and show the class the accumulation of iron particles on the magnetic stir bar.
NOTES AND REFERENCES
1. The iron or steel wire must be cleaned prior to use to remove any protective coating of oil or grease. This may be done by wiping the wire with a paper towel and a solvent such as acetone, ethanol, paint thinner or turpentine. Steel wire is frequently lightly galvanized, and this will result in a brief, rapid evolution of H2 gas while the zinc coating dissolves, followed by a slower evolution of gas as the iron itself dissolves.
2. Sodium or ammonium thiocyanate may be used in place of potassium thiocyanate. Ferric ammonium sulfate may be used in place of ferric chloride.
3. White blotter paper or water color paper work best for this purpose. Typically, we paint a US flag on the paper, but the picture can be anything desired, or it may be a message such as Happy Memorial Day Weekend, Happy Birthday, or even a reminder of an upcoming exam. See Chen, Philip S. Entertaining and Educational Chemical Demonstrations; Chemical Elements Publishing Co.: Camarillo, CA, 1974; pp 20.
4. The solution in the bottle or flask contains 5 grams ferric ammonium sulfate in 500 mL of water. Each of the beakers contains about 0.5 g of the solid reagent dissolved in 3 to 5 mL of water: (1) potassium thiocyanate, (2) barium chloride, (3) potassium ferrocyanide, (4) tannic acid, (5) sodium bisulfite (NaHSO3), (6) tartaric acid. Note that the white color produced in (2) is a precipitate of BaSO4 and is not due to iron chemistry. The colors produced are (1) red, (2) white, (3) blue, (4) black, (5) orange, (6) yellow. See Ford, Leonard A. Chemical Magic, 2nd ed.; Dover: Mineola, NY, 1993; pp 4.
5. Aluminum is not used for this purpose because it rapidly becomes coated with a dense, tough, and insulating coating of aluminum oxide and therefore does not bond well to iron. 6. See (a) Wright, Stephen W.; Folger, Marsha R.; Quinn, Ryan P., J. Chem. Ed. 2005, 82, 1633; (b) Wright, Stephen W., J. Chem. Ed. 2005, 82, 1648A – 1648B.
For Laboratory Work: Please refer to the ACS Guidelines for Chemical Laboratory Safety in Secondary Schools (2016).
For Demonstrations: Please refer to the ACS Division of Chemical Education Safety Guidelines for Chemical Demonstrations.
Other Safety resources
RAMP: Recognize hazards; Assess the risks of hazards; Minimize the risks of hazards; Prepare for emergencies