Materials Included In Kit

Additional Materials Required

Safety Precautions

Potassium ferricyanide, potassium ferrocyanide and potassium thiocyanate are dangerous if heated or in contact with concentrated acids since toxic hydrogen cyanide gas may be liberated. Potassium thiocyanate is moderately toxic by ingestion. Potassium ferricyanide, potassium ferrocyanide and ferrous sulfate are slightly toxic by ingestion. Iron(II) sulfate is corrosive to skin, eyes and mucous membranes. Iron(III) chloride and iron(III) nitrate are corrosive and may be skin and tissue irritants. Avoid body contact with all chemicals. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron or laboratory coat. Wash hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information. Please follow all laboratory safety guidelines.

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 the solutions in the test tubes down the drain with excess water according to Flinn Suggested Disposal Method #26b. Flush any excess potassium thiocyanate, iron(II) sulfate, or iron(III) chloride solution down the drain according to Flinn Suggested Disposal Method #26b. Dispose of any excess ferricyanide and ferrocyanide solution according to Flinn Suggested Disposal Method #14. Dispose of any excess iron(III) nitrate solution according to Flinn Suggested Disposal Method #24b.

Lab Hints

• Enough materials are provided in this kit for 30 students working in pairs or for 15 groups of students. Both parts of this laboratory activity can reasonably be completed in one 50-minute class period. The prelaboratory assignment may be completed before coming to lab, and the worksheet may be completed the day after the lab.
• Instead of using the iron(III) nitrate solution provided in the kit for the unknown, you can have the students test any iron solution in your chemical stockroom.

Answers to Prelab Questions

1. 1. Click on the Reduction Potential virtual activity. Follow the on-screen prompts to observe the reactions as various metals are placed in different metal-containing solutions. Observe how the reactions correlate with the standard reduction potential chart.
   2. Record the ID number of your unknown. Place the unknown metal into the different flasks to determine the identity based on its reactivity and the information you learned in Part a.

      The unknown will be randomly selected from a collection of 4 metals. Each unknown will have an ID number that is also randomized to prevent students from sharing answers. If the ID number ends in a one, the unknown is lead. If the ID number ends in a two, the unknown is nickel. If the ID number ends in a three, the unknown is chromium. If the ID number ends in a four, the unknown is manganese.

2. Determine the oxidation number of the iron atom in each solution provided in this lab, and record this information on your data sheet.

Show all work.

FeSO₄: The sulfate ion has a 2– charge; iron must be 2+.
FeCl₃: The chlorine ion has a 1– charge and there are three of them; iron must be 3+.
K₄[Fe(CN)₆]: Potassium is 1+ multiplied by 4, the cyanide ion is 1– multiplied by 6; iron must be 2+.
K₃[Fe(CN)₆]: Potassium is 1+ multiplied by 3, the cyanide ion is 1– multiplied by 6; iron must be 3+.
KSCN: No iron is present.

Sample Data

Redox Observations

Determination of Unknown

Answers to Questions

1. How can potassium thiocyanate be used to confirm that Fe²⁺ ions have been oxidized to Fe³⁺? Will tube 1 eventually turn Prussian blue?

A solution of Fe³⁺ ions turns dark red when potassium thiocyanate is added. As Fe²⁺ is oxidized by atmospheric oxygen to form Fe³⁺ and will take on the Prussian blue color.

2. How can potassium ferricyanide be used to confirm that Fe³⁺ ions have been reduced to Fe²⁺?

A solution of Fe²⁺ turns dark blue when potassium ferricyanide is added.

3. Use the oxidation state rules to assign oxidation states for the indicated atoms in each oxidizing agent and its product. Show your work.

MnO₄^–: The oxygen ion has a 2– charge, multiplied by 4 gives 8–; manganese must be 7+.
H₂O₂: The hydrogen ion has a 1+ charge, multiplied by 2 gives +2; each oxygen must be 1–.
OCl^–: The oxygen ion has a 2– charge, and chlorine is almost always 1–.
IO₃^–: The oxygen ion has a 2– charge, multiplied by 3 gives 6–; iodine must be 5+.
SO₄^2–: The oxygen ion has a 2– charge, multiplied by 4 gives 8–; sulfur must be 6+.

4. Circle the correct choices to complete the following definitions.

1. An oxidizing agent is a substance that causes the (oxidation/reduction) of another reactant in a redox reaction. The oxidation state of the oxidizing agent (increases/decreases), and the oxidizing agent itself undergoes (oxidation/reduction) during the reaction.
2. A reducing agent is a substance that causes the (oxidation/reduction) of another reactant in a redox reaction. The oxidation state of the reducing agent (increases/decreases), and the reducing agent itself undergoes (oxidation/reduction) during the reaction.

References

Bilash, B.; Gross, G.; Koob, J. A Demo A Day™—Another Year of Chemical Demonstrations, Vol 2; Flinn Scientific: Batavia, IL, 1998; pp 244–246.

Tzimopoulos, N. D.; Metcalfe, H. C.; Williams, J. E.; Castka, J. F. Modern Chemistry Laboratory Experiments; Holt, Rinehart and Winston: New York, 1990; p 63.

Introduction

Iron exists in the body in two forms—iron(II), Fe²⁺, and iron(III), Fe³⁺, ions. Both forms of iron are important in the absorption, storage and utilization of iron by the body. It is easy to distinguish between solutions of iron(II) and iron(III) ions by performing redox reactions between the two oxidation states.

Concepts

• Complex ions
• Oxidation numbers
• Oxidation–reduction
• Transition metals
• Reduction potentials

Background

Oxidation–reduction reactions are a major class of chemical reactions. Often called redox for short, these reactions are defined as any reaction in which electrons are transferred from one substance to another. Oxidation occurs when a substance loses electrons and reduction occurs when a substance gains electrons. Due to this loss and gain of electrons, oxidation and reduction always occur in tandem—one cannot occur without the other. Because of this paired relationship, substances that are used to cause oxidation or reduction are called oxidizing and reducing agents, respectively. The substance that accepts electrons in a redox reaction is called the oxidizing agent—by accepting electrons, it causes the oxidation of another substance. Similarly, the substance that loses electrons in a redox reaction is called the reducing agent because it causes the reduction of another substance.

The loss and gain of electrons by the reactants in a chemical reaction is not always obvious from the formulas of the reactants and products. A method based on oxidation states has been developed to identify oxidation–reduction reactions, to determine whether a substance has been oxidized or reduced and to count the electrons that are lost or gained as a result.

The following rules are used to assign oxidation states:

1. All pure elements have an oxidation number of 0: Mg in Mg(s) has a 0 oxidation number. Each O in O₂(g) is also 0.
2. Monatomic ions have oxidation numbers that match the charge of the ion: Na+ = +1 or F– = –1.
3. In a neutral compound, the sum of all oxidation numbers is 0: NaCl = (+1) + (–1) = 0 or AlPO₄ = (+3) + (+5) + 4(–2) = 0.
4. In a polyatomic ion, the sum of all oxidation numbers equals the charge of the ion. For example, CO₃^2–: C = +4 and O = –2: (+4) + 3(–2) = –2 = charge.
5. In a binary compound, the more electronegative element has an oxidation number equal to the charge it would have as a monatomic ion.
6. Hydrogen typically has an oxidation number of +1, unless it is attached to a metal, then it is –1. For example, HCl (H = +1, Cl = –1) and NaH (Na = +1, H = –1).
7. Group I and II elements have oxidation numbers of +1 and +2, respectively. For example, NaF (Na = +1, F = –1) and CaCl₂ (Ca = +2, Cl = –1).
8. Aluminum has an oxidation number of +3. For example, AlCl₃ (Al = +3, Cl = –1).
9. Fluorine has an oxidation number of –1 in its compounds.
10. Oxygen typically has an oxidation number of –2 (CaO; Ca = +2, O = –2). When oxygen is combined with fluorine, it can have a +1 or +2 oxidation number (OF₂; O = +2, F = –1). When oxygen is part of a peroxide, it has a –1 oxidation number (H₂O₂; H = +1, O = –1).

A reaction is classified as a redox reaction if the oxidation states of the reactants change. Oxidation is an increase in oxidation state (equivalent to a loss of electrons). Reduction is a decrease in oxidation state (equivalent to a gain of electrons). A simple way to remember the difference is “LEO says GER” (see Figure 1).

Consider the reaction of Fe²⁺ ions with chlorine (Equation 1). Iron is oxidized—the oxidation state of iron increases from +2 to +3. Chlorine is reduced—the oxidation state of chlorine decreases from 0 to –1.

For every redox reaction, the two separate half-reactions can be written. The oxidation half-reaction shows the substance that is oxidized, the product resulting from oxidation and the number of electrons lost in the process. The number of electrons lost is equal to the difference in oxidation states between the reactant and product. The reduction half-reaction shows the substance that is reduced, the number of electrons gained in the process and the product resulting from the reduction. The oxidation and reduction half-reactions for the redox reaction of Fe²⁺ with chlorine are shown. The oxidation half-reaction must be multiplied by a factor of two so the number of electrons lost by Fe²⁺ will be equal to the number of electrons gained by chlorine.

Fe²⁺ → Fe³⁺ + e^– Oxidation half-reaction
Cl₂ + 2e^– → 2Cl^– Reduction half-reaction

A common type of redox reaction is a single replacement reaction, which typically involves metals. The ability to replace another metal determines a metal’s reactivity—the better its ability to replace another metal, the more reactive the metal is. The activity series of metals is a list that places the metals in order of reactivity (see Figure 2).

The metals at the bottom are more reactive and can replace the metals above them. Reactivity increases as you move down the list.

Many transition metals exhibit the ability to exist as relatively stable ions in different oxidation states. As discussed earlier, iron can be found as the Fe²⁺ ion or the Fe³⁺ ion depending on the compound. The variable valence states can be explained by looking at the electron configuration of iron, which is [Ar]4s²3d⁶. When transition metal atoms form positive ions, the outer s electrons are lost first because the inner d sublevels are lower in energy (more stable) than the outer s sublevels. In the iron(II) ion, the two 4s electrons have been lost, leaving [Ar]3d⁶. In the iron(III) ion, the two 4s elections and one 3d electron have been removed, leaving [Ar]3d⁵. The iron(III) ion is more stable than the iron(II) ion since its d orbital is half-filled. Half-filled orbitals (and filled orbitals) have been shown to have greater stability. Therefore, a compound or a solution containing the iron(II) ion will slowly oxidize to the iron(III) state on exposure to air due to the greater stability of the Fe³⁺ ion.

In order to distinguish between iron(II) and iron(III) ions, potassium ferrocyanide, K₄[Fe(CN)₆]•3H₂O, and potassium ferricyanide, K₃[Fe(CN)₆], complexes are used in this experiment. The cyano group in each complex has a charge of –1, and potassium has a charge of +1. This means the complex ferrocyanide, [Fe(CN)₆]^4–, contains iron in the +2 oxidation state while the complex ferricyanide, [Fe(CN)₆]^3–, contains iron in the +3 oxidation state. A deep-blue (Prussian blue) precipitate results when either complex ion combines with iron in a different oxidation state from that present in the complex. This provides a means of identifying either iron ion. When a solution of iron(II) is mixed with ferricyanide, iron(III), a deep-blue precipitate is formed; likewise, when a solution of iron(III) is mixed with ferrocyanide, iron(II), the same deep-blue precipitate is observed.

The deep-blue precipitate, Prussian blue, has the composition of Fe₄[Fe(CN)₆]₃. Prussian blue has been used as a pigment in printing inks, paints, cosmetics (eye shadow), artist colors, carbon paper and typewriter ribbons. The thiocyanate ion, SCN^–, provides an excellent confirming test for the Fe³⁺ ion. The water soluble, blood-red [Fe(SCN)₃•(H₂O)₅]⁺ complex is formed from the Fe³⁺ ion while no complex is formed with the Fe²⁺ ion.

Experiment Overview

The purpose of this experiment is to investigate the reactions of Fe²⁺ and Fe³⁺ ions by adding various complex ions to solutions of iron(II) or iron(III)—observe the formation of the beautifully colored Prussian blue precipitate or of the deep blood-red complex, the confirming test for iron(III).

Materials

Prelab Questions

1. 1. Click on the Reduction Potential virtual activity. Follow the on-screen prompts to observe the reactions as various metals are placed in different metal-containing solutions. Observe how the reactions correlate with the standard reduction potential chart.
   2. Record the ID number of your unknown. Place the unknown metal into the different flasks to determine the identity based on its reactivity and the information you learned in Part a.

Unknown ID: _______
Metal identity: _____________

2. Determine the oxidation number of the iron atom in each solution provided in this lab, and record this information on your data sheet. Show how you determined the oxidation number.

Safety Precautions

Potassium ferricyanide, potassium ferrocyanide and potassium thiocyanate are dangerous if heated or in contact with concentrated acids since toxic hydrogen cyanide gas may be liberated. Potassium thiocyanate is moderately toxic by ingestion. Potassium ferricyanide, potassium ferrocyanide and ferrous sulfate are slightly toxic by ingestion. Iron(II) sulfate is corrosive to skin, eyes and mucous membranes. Iron(III) chloride and iron(III) nitrate are corrosive and may be skin and tissue irritants. Avoid body contact with all chemicals. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron or laboratory coat. Wash hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety. Please follow all laboratory safety guidelines.

Procedure

1. Place six test tubes in a test tube rack. Label tubes 1–3 as Fe²⁺ and tubes 4–6 as Fe³⁺.
2. Add approximately 10 mL of 0.02 M iron(II) sulfate solution and 10 mL of distilled or deionized water to test tubes 1–3. Stopper the tubes and invert to mix.
3. Add approximately 10 mL of 0.02 M iron (III) chloride solution and 10 mL of distilled or deionized water to test tubes 4–6. Stopper the tubes and invert to mix.

Part A. Ferrocyanide Ions, [Fe(CN)₆]^4– (Iron in the +2 Oxidation State)

4. Add 2–3 drops of 0.1 M potassium ferrocyanide solution to tube 1. Record observations in the data table.
5. Add 2–3 drops of 0.1 M potassium ferrocyanide solution to tube 4. Record observations in the data table.

Part B. Ferricyanide Ions, [Fe(CN)₆]^3– (Iron in the +3 Oxidation State)

6. Add 2–3 drops of 0.1 M potassium ferricyanide solution to tube 2. Record observations in the data table.
7. Add 2–3 drops of 0.1 M potassium ferricyanide solution to tube 5. Record observations in the data table.

Part C. Thiocyanate Ions, SCN^–

8. Add 2–3 drops of 0.1 M potassium thiocyanate solution to tube 3. Record observations in the data table.
9. Add 2–3 drops of 0.1 M potassium thiocyanate solution to tube 6. Record observations in the data table.
10. Dispose of the solutions in the test tubes as directed by your teacher.

Part D. Determination of Unknown

11. Label three test tubes 1–3, and obtain 30 mL of an unknown iron solution from your teacher.
12. Add 10 mL to each of three test tubes along with 5 mL of distilled or deionized water. Stopper the tubes and invert to mix.
13. Add 1–2 drops of 0.1 M potassium ferrocyanide solution to tube 1. Record observations in the data table.
14. Add 1–2 drops of 0.1 M potassium ferricyanide solution to tube 2. Record observations in the data table.
15. Add 1–2 drops of 0.1 M potassium thiocyanate solution to tube 3. Record observations in the data table.
16. Based on your observations, determine whether your unknown contains iron in the +2 or +3 oxidation state.

Student Worksheet PDF

14160_Student1.pdf