Materials Included In Kit

Additional Materials Required

Safety Precautions

Potassium ferricyanide solution is a mild skin irritant. Contact with strong acids may liberate toxic hydrogen cyanide gas; avoid contact with strong acids. Phenolphthalein indicator solution is an alcohol-based solution; it is a flammable liquid and is moderately toxic by ingestion. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. Dispose of all Petri dishes (each still containing the nails, agar and indicator solutions) in the trash according to Flinn Suggested Disposal Method #26a.

Teacher Tips

• To bend the nails, hold each nail with one pair of pliers just above where the bend is to be. Then, use the other pair of pliers to bend the other half of the nail down. The bend does not need to be exactly 90°. To save time, this may be done before the class period.
• When the 0.1 M potassium ferricyanide and phenolphthalein indicator solutions are added to the agar suspension, the resulting color should be yellow. If it is not, have students add additional potassium ferricyanide dropwise until the suspension becomes yellow in color.
• Observing the plates after 24 hours is optimal. After a few days, the nails will begin to rust and mold will most likely be present.
• This procedure may be used as a demonstration. Simply prepare the plates as directed in this procedure, then display them on an overhead projector.
• The plates may be prepared prior to lab if desired. To do this, simply dissolve about 15 g of powdered agar in 1500 mL of boiling water. Add about 3 mL each of the 0.1 M potassium ferricyanide and phenolphthalein indicator solutions to the agar suspension. The resulting solution should be yellow in color. If it is not, add additional potassium ferricyanide until the solution is yellow in color. Pour the agar suspension into 30 plates and allow them to cool and harden. In lab, the students can simply press the nails into the agar, then view the results at the end of lab and during the next lab period.

Sample Data

Observations

Answers to Questions

1. What is the purpose of cleaning the iron nails with steel wool before performing the laboratory procedure?

   There may have been a metal oxide coating on the nails, which could interfere with the desired redox reactions.

2. What effect does bending the iron nail have on the oxidation process?

   Bending the nail puts stress on the nail. Oxidation occurs most readily at the points on the nail which have been stressed in the manufacturing process (the point and head of the nail).

3. Which parts of the bent and control iron nails are oxidized most readily? What evidence supports this? Propose an explanation for why this occurs.

   Oxidation (the blue color) occurs at the sharp points or turns in the nail. The reason why this occurs is below. Because the answer involves fairly high level concepts, allow for creativity in student answers.
   The iron atoms in the regions of the nail containing turns or sharp points have been stressed in the manufacturing process. Stress tends to promote oxidation by disrupting the matrix of iron atoms and creating defects on an atomic level. It is easier for the iron atoms at the sharp points or turns to leave the surface of the nail and dissolve in the surrounding solution. To dissolve in solution, however, the iron atom must be oxidized to form an ion. For iron to be oxidized, there must be some oxidizing agent (oxygen) present which will accept the electrons that the iron atoms are donating.
   Although the sharp points or turns are where the iron atoms leave the nail and go into solution as iron ions, these are not the points where the electrons that are left over from the oxidation of iron go into solution. Instead, the electrons travel down the network of metal atoms to the straight parts of the nail. Here, they participate in the reduction of oxygen, adding OH^– ions to solution near the straight parts of the nail.
   So, near the turns or sharp points of the nail, the agar suspension contains Fe²⁺ ions, while the area surrounding the straight parts of the nail contains OH^– ions. The Fe²⁺ ions react with the potassium ferricyanide in suspension to give ferrous ferricyanide. Therefore, near the turns or sharp points of the nail, the solution turns blue. The presence of OH^– ions in solution creates a basic environment. In base, the indicator phenolphthalein turns pink. Therefore, near the straight parts of the nail, the solution turns pink.

4. Does zinc protect the iron nail from oxidation? What evidence supports this?

   Yes. There is no blue color around the zinc-coated nail. Zinc is above iron in the activity series, so it should protect iron from oxidation.

5. Does copper protect the iron nail from oxidation? What evidence supports this?

   No. Blue appears around the copper-coated nail. Copper is below iron in the activity series, so it cannot protect iron from oxidation. In fact, the iron serves as a protector to the copper in this case.

6. Predict another metal which will protect an iron nail. Why should this metal serve as a protector to the iron nail?

   Student answers will vary. Valid answers include any of the metals that appear above iron in the activity series (Table 1) of the Background Section; however, students should consider the properties of the metal also. For instance, none of the alkali metals would be good choices due to their violent reactivity in water.

References

We thank Greg Kifer and Rhonda Rheist of Olathe North High School, Olathe, KS, for providing us with tips and background information.

Introduction

The corrosion of iron, better known as rusting, is an oxidation–reduction process that destroys iron objects left out in open, moist air. Can corrosion be prevented or slowed down? In this laboratory activity, the corrosion of several iron nails subjected to different conditions will be studied to determine which conditions promote or prevent the corrosion of iron.

Concepts

• Redox reactions
• Corrosion
• Activity series

Background

Redox Reactions

Oxidation–reduction, or redox, reactions are reactions in which electrons are transferred from one element to another. There are two key parts present in every redox reaction—an element that is oxidized and an element that is reduced. Oxidation of an element occurs when the element donates electrons. The net result is that the charge on an oxidized substance is increased during the chemical reaction. Reduction of an element occurs when an element accepts electrons. As a result, the charge on a reduced element is decreased during a chemical reaction.

Examine Equation 1. It is a redox reaction. In this reaction, electrons are transferred from aluminum atoms to cupric ions. Aluminum donates electrons and its charge increases—it is oxidized. The cupric ion accepts electrons and its charge decreases—it is reduced.

Another way to look at this reaction is in terms of oxidizing and reducing agents. The substance that accepts electrons (and is thus reduced) in a chemical reaction is the oxidizing agent while the substance that donates electrons (and is thus oxidized) is the reducing agent. In Equation 1, the cupric ion accepts electrons, so it is the oxidizing agent. In other words, the cupric ion oxidizes aluminum. Conversely, aluminum donates electrons, which makes it the reducing agent—it reduces the cupric ion.

An easy way to identify the oxidizing and reducing agents in a chemical reaction is to look at the difference in charge on an element between the reactants and the products. The charge on the oxidized species increases during a chemical reaction, while the charge on the reduced species decreases during a chemical reaction. Because it is the reduced substance’s charge that decreases, this is often a convenient way to remember which species is reduced and is, therefore, the oxidizing agent.

The Rusting Process

When iron metal is exposed to oxygen and water, a familiar result is observed—rust. In the rusting process, which consists of several steps, iron is first oxidized to the ferrous ion, Fe²⁺, while oxygen is reduced according to Equation 2. Breaking this reaction up into its component half reactions gives Equations 3 and 4. Both of these half reactions must occur simultaneously when iron is exposed to oxygen and water. Because the products of these half reactions, Fe²⁺ and OH^–, are in contact with each other, they react together to form solid ferrous hydroxide, Fe(OH)₂, according to Equation 5. The ferrous hydroxide is still in contact with oxygen and water, and it reacts with them to produce rust, or ferric oxide, Fe₂O₃, according to Equation 6. As you know from practical experience, iron does not rust immediately after being exposed to oxygen and water. Instead, the rusting process takes time. In this laboratory activity, the initial steps (Equations 2–4) in the rusting process will be observed over a 24-hour period.

The Activity Series

The activity series of metals is a scheme that places the metals in order of reactivity (see Table 1). Reactivity can be defined as the ease of oxidation. The metals at the top are the most reactive, or the most easily oxidized, and the reactivity of the metals decreases as you go down the list. This list can be used to determine which metals can serve as protectors to other metals from oxidation by oxygen and water. Each of the metals listed above oxygen in Table 1 can be oxidized by oxygen in the presence of water. When oxidized, these metals form their metal oxides. Familiar metal oxides include the tarnish that forms on the surface of copper, or the rust that plagues iron. Because conversion of a metal to its metal oxide is many times a costly, unwanted process, it is important to find ways to prevent or slow down this process called corrosion.

When two metals are present together and both are exposed to oxygen and water, the more reactive, or most easily oxidized, metal will be oxidized to its metal oxide by the oxygen first. The less reactive metal will not be affected by the presence of the oxygen until the more reactive metal is completely oxidized. In effect, the more reactive metal protects the less reactive metal from the oxidizing effects of oxygen.

For example, if manganese and zinc were present together and both were exposed to oxygen and water, manganese would be oxidized by oxygen to form manganese oxide. The zinc would not initially be affected by the presence of the oxygen. Not until all of the manganese had been converted to manganese oxide would the zinc feel the oxidizing effects of the oxygen and be converted to zinc oxide. In this example, the manganese serves as a protector to the zinc—the manganese prevents the zinc from being oxidized immediately.

Indicator Reactions

In this laboratory activity, two indicators will be used to identify where oxidation and reduction occur on a iron nail. The oxidation and reduction reactions are given in Equations 3 and 4. The products of these reactions, Fe²⁺ and OH^–, will react differently with two indicators, potassium ferricyanide and phenolphthalein. Fe²⁺ ions can react with potassium ferricyanide to produce ferrous ferricyanide, which is also called Turnbull’s blue, according to Equation 7. Thus, formation of a blue precipitate indicates that oxidation (see Equation 3) has occurred. In the presence of OH^– ions, phenolphthalein turns pink. Thus, the appearance of a pink color indicates that reduction has occurred (see Equation 8).

Materials

Safety Precautions

Potassium ferricyanide solution is a mild skin irritant. Contact with strong acids may liberate toxic hydrogen cyanide gas; avoid contact with strong acids. Phenolphthalein indicator solution is an alcohol-based solution; it is a flammable liquid and is moderately toxic by ingestion. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

1. Pour 100 mL of distilled or deionized water into a 250-mL beaker. Heat the water to boiling on a hot plate.
2. While the water is coming to a boil, clean four iron nails with steel wool. Bend one of the nails using two pairs of pliers so that it has a 90° bend in it. Cut a 4-inch piece of copper wire using scissors. Tightly wrap one of the straight nails with a 4-inch piece of copper wire. Tightly wrap another straight nail with a 5-inch strip of zinc metal. See Observations section below. The fourth nail, to which nothing has been done, is the control.
3. Once the water has come to a rapid boil, add 1 g of powdered agar with continued stirring. Stir until the agar forms a uniform suspension. Be careful not to burn the agar. Remove the agar from the hot plate. Turn off the hot plate.
4. Add 5 drops of potassium ferricyanide solution and 5 drops of phenolphthalein solution to the agar suspension. Stir to mix the suspension thoroughly. The suspension should be yellow in color.
5. Allow the agar to cool, but not harden.
6. Place the control nail and the bent nail in one of the Petri dish bottoms. Make sure these nails are not touching.
7. Place the two metal-wrapped nails in the second Petri dish bottom. Make sure these nails are not touching.
8. Pour the agar suspension into both Petri dishes. Fill the dishes so that the nails are completely submerged.
9. Cover both Petri dishes and set them aside for viewing later.
10. Observe the Petri dishes at the end of the class period. Record your observations in the drawing below by indicating where any color changes may be observed.
11. Observe the Petri dishes the next day. Record your observations in the drawing below by indicating where any color changes may be observed.

Student Worksheet PDF

11871_Student1.pdf