Materials Included In Kit

Additional Materials Required

Prelab Preparation

Iron(II) sulfate solution, 0.2 M: Dissolve 5.6 g of iron(II) sulfate heptahydrate (FeSO₄•7H₂O) in about 50 mL of distilled or deionized water. Stir to dissolve and dilute to 100 mL with water. Note: Prepare this solution fresh the day of use.

Safety Precautions

Silver nitrate is slightly toxic by ingestion and will stain skin and clothing. Copper(II) sulfate and iron(II) sulfate solutions are toxic by ingestion. Magnesium nitrate solution is a skin and eye irritant. Metal pieces may have sharp edges—handle with care. Avoid contact of all chemicals with eyes and skin. 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. Remind students to wash their hands thoroughly with soap and water before leaving the lab.

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. The contents of the reaction plate may be safely discarded in the solid waste according to Flinn Suggested Disposal Method #26a. Excess iron(II) sulfate solution may be disposed of down the drain with plenty of excess water according to Flinn Suggested Disposal Method #26b. Save the other solutions in properly labeled bottles for future use.

Lab Hints

• Enough materials are provided in this kit for 30 students working in pairs or for 15 groups of students. The laboratory work for this microscale experiment can easily be completed in a typical 50-minute lab period. The Prelab Questions may be assigned as homework in preparation for lab. To facilitate the lab write-up, review the Prelab Questions during the lab period.
• Metal ribbon, sheet, turnings, foil, etc. may all be used in this experiment. (Wire is not recommended.) Sometimes the only evidence that a reaction has occurred is a slight darkening or coating of the metal. The best results were obtained using small metal strips cut from thin metal sheets or foil—the metal pieces are easy to remove and examine and any changes are readily apparent.
• Based on the use of net ionic equations to describe single replacement reactions, the choice of the anion in the metal salt solution should not affect whether a metal reacts with a particular metal salt solution. This is not always the case, however. In some reactions, the anion does matter. For example, aluminum metal will react with copper(II) chloride, but not with copper(II) sulfate. The properties of the anion influence both the acidity of the solution and the permeability of the metal surface.
• Sanding and cutting metal pieces is time-consuming. We recommend that the teacher sand large pieces of metal ribbon, sheet or foil before lab. Students may then cut the metal into 1-cm² pieces as needed. This is more efficient than having students sand tiny strips of metal. With the exception of iron strips, all of the recommended forms of the metals may be cut using regular scissors. Use heavy-duty scissors (Flinn Catalog No. AP8949) to cut iron strips.
• Place metal ion solutions in dropper bottles or in labeled pipet sets for student use. When the lab is over, empty and rinse the dropper bottles or pipets and save them for use in subsequent years. A filled pipet (2–3 mL) should be enough for two groups of students.
• Introduce the experiment by demonstrating the reaction of aluminum foil with copper(II) chloride, as described in the Background section—see the Foiled Again! chemical demonstra tion kit available from Flinn Scientific (Catalog No. AP5936). An alternative example of the relationship between single replacement reactions and metal activity is the reaction of mossy zinc with tin(II) chloride—see the Floating Tin Sponge chemical demonstration kit available from Flinn Scientific (Catalog No. AP4425).
• Other metals and metal ion solutions may be added to this experiment to build a more inclusive list of metal activity—lead and lead(II) nitrate, tin and tin(II) chloride, manganese and manganese(II) chloride, nickel and nickel(II) nitrate, etc. Use of heavy metal ion solutions will require dedicated heavy metal disposal. Although it is a very reactive metal, aluminum is often omitted from metal activity experiments because it gives anomalous results. Aluminum is a “self-protecting” metal— the aluminum oxide coating is extremely tough, durable, and essentially unreactive. Strong acid is needed to remove the oxide coating on aluminum.
• Spray-n-Wash^® laundry product will remove silver nitrate stains from clothing.
• Consider removing the metal piece from one of the “slow” reactions in the sample data table and examining the reaction under a microscope. You should see pitting or scarring of the metal surface. Ask your colleagues in the Biology department to lend you a microscope (in exchange perhaps for a future solution prep).

Teacher Tips

• Differences in metal activity lead to galvanic corrosion, which occurs when two dissimilar metals are brought together by means of an electrolyte. (The metals do not actually have to be touching each other. The electrolyte acts as a conductor that connects the two metals as part of an electrochemical cell.) When two dissimilar metals are coupled, corrosion of the more active metal will accelerate, while corrosion of the less active metal will slow down or even stop. One of the most famous examples of galvanic corrosion involves the damage to the Statue of Liberty. Between 1984–1986, the Statue of Liberty underwent a two-year, $90 million dollar restoration to repair and replace the interior iron scaffolds supporting the copper “skin” of the monument itself. The iron skeleton was initially separated from the copper skin by means of a shellac-coated insulator. One hundred years of exposure to salt and humidity broke the insulating seal between the iron and copper. This resulted in extensive damage to the iron framework, since iron is more active than copper.

Answers to Prelab Questions

1. According to the Background information, magnesium is more active than aluminum. Predict which combination of metal and metal ion solution will show evidence of a chemical reaction: (a) Mg(s) and Al₂(SO₄)₃(aq) or (b) Al(s) and MgSO₄(aq).
   1. Magnesium metal will react with aluminum ions from aluminum sulfate. The magnesium metal will become coated and eventually disintegrate as new aluminum particles break off from the metal surface.
   2. Aluminum metal will not react with magnesium ions from magnesium sulfate.
2. Write the balanced chemical equation for the single replacement reaction that is expected to occur (1a or 1b).

   3Mg(s) + Al₂(SO₄)₃(aq) → 3MgSO₄(aq) + 2Al(s)

3. Which substance is being oxidized in this reaction? Which substance is being reduced?

   Magnesium is being oxidized; aluminum ions in aluminum sulfate are being reduced.

4. Write the oxidation and reduction half-reactions for the overall reaction.

   Mg(s) → Mg²⁺(aq) + 2e^– Oxidation half-reaction
   Al³⁺(aq) + 3e^– → Al(s) Reduction half-reaction

Sample Data

Answers to Questions

1. Which metals reacted with (a) the most metal ion solutions and (b) the fewest metal ion solutions?

   Magnesium reacted with all of the metal ion solutions. Copper reacted with only one of the metal ion solutions.

2. Compare the general trend in the reactivity of Cu, Fe, Mg and Zn and rank the metals from most active (first) to least active (last).

   The general trend is that each metal reacted with a different number of metal ion solutions. Thus, the most active metal reacted with four metal ions, while the least active metal reacted with only one metal ion.
   Mg > Zn > Fe > Cu

3. Because silver metal is expensive, it was not used in this experiment. Based on the observed reactions of Cu, Fe, Mg and Zn with silver nitrate, explain why it was not necessary to test silver in order to determine its activity.

   All of the metals tested reacted with silver nitrate. Single replacement reactions occur in one direction only, with the more active metal replacing the less active metal ion. Therefore, all of the metals tested (Cu, Fe, Mg and Zn) must be more active than silver.

4. Rewrite the activity series of the metals (Question 3) to include silver.

   Mg > Zn > Fe > Cu > Ag

5. Write the balanced, net ionic equation for each single replacement reaction of a metal (M) with a metal ion (N^x+) observed in this activity. Hint: Substitute the symbols of the metals and metal ions into the following equation and remember to balance the charges.

   M(s) + N^x+(aq) → M^y+(aq) + N(s)
   Cu(s) + 2Ag⁺(aq) → Cu²⁺(aq) + 2Ag(s)
   Fe(s) + Cu²⁺(aq) → Fe²⁺(aq) + Cu(s)
   Fe(s) + 2Ag⁺(aq) → Fe²⁺(aq) + 2Ag(s)
   Mg(s) + Cu²⁺(aq) → Mg²⁺(aq) + Cu(s)
   Mg(s) + Fe²⁺(aq) → Mg²⁺(aq) + Fe(s)
   Mg(s) + 2Ag⁺(aq) → Mg²⁺(aq) + 2Ag(s)
   Mg(s) + Zn²⁺(aq) → Mg²⁺(aq) + Zn(s)
   Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
   Zn(s) + Fe²⁺(aq) → Zn²⁺(aq) + Fe(s)
   Zn(s) + 2Ag⁺(aq) → Zn²⁺(aq) + 2Ag(s)

6. The substance that accepts electrons and is reduced in an oxidation–reduction reaction is called the oxidizing agent (it causes the oxidation of another substance). The substance that gives up electrons and is oxidized acts as a reducing agent (it causes the reduction of another substance). An active metal is a good reducing agent: True or False? Explain.

   An active metal is a good reducing agent: True. Magnesium was the most active metal tested. It reduced four different metal ions.

References

This activity was adapted from Chemical Reactions, Vol. 6 in the Flinn ChemTopic™ Labs series; Cesa, I., Editor; Flinn Scientific: Batavia, IL (2004).

Introduction

The usefulness of metals in structural and other applications depends on their physical and chemical properties. Although iron is the most common metal used in manufacturing, it must be protected against corrosion because iron rusts easily. Copper is used in electrical wiring because it conducts electricity extremely well and resists corrosion better than many metals. Gold is a highly valuable jewelry metal because it is essentially unreactive. How can we determine the relative reactivity of different metals?

Concepts

• Activity series
• Single replacement reactions
• Oxidation–reduction
• Net ionic equations

Background

A ranking of metals in order of their tendency to react with acids and water is called an activity series. Comparing the reactions of sodium, magnesium, and aluminum reveals that sodium reacts violently with acids and water, magnesium reacts with acids and hot water, and aluminum reacts only with acids. Based on this trend in reactivity, sodium is more active than magnesium, which is more active than aluminum. The activity series for these metals would be written as Na > Mg > Al.

Another way to determine the activity of metals is to compare the reactions of metals with different metal ions. Consider, for example, what happens when a piece of aluminum foil is placed in a solution of copper(II) chloride. A vigorous reaction is observed—heat is released, the blue color due to copper(II) ions fades, the aluminum foil disintegrates, and a new, reddish brown solid appears in the reaction mixture. The reaction is summarized in Equation 1. Copper(II) ions from copper(II) chloride are converted to copper metal, and aluminum metal is converted to aluminum cations in aluminum chloride, which is soluble in water. In contrast, when a piece of copper metal is placed in a solution of aluminum chloride, no reaction takes place (Equation 2).

The reaction of aluminum with copper(II) chloride is classified as a single replacement reaction—aluminum reacts with and “replaces” copper ions in copper(II) chloride. Single replacement reactions will occur spontaneously in one direction only (compare Equations 1 and 2). A more active metal always replaces the ion of a less active metal in a compound. In general, the activity of a metal may be defined as follows: An active metal will react with a compound of a less active metal, which is converted to its “free element” form. The more active metal forms a new compound containing metal cations. Based on Equation 1, aluminum is more active than copper.

Single replacement reactions are examples of oxidation–reduction reactions. Oxidation is defined as the process of losing electrons, and a substance that loses electrons during a chemical reaction is said to be oxidized. If one substance loses electrons during the course of a chemical reaction, another substance must gain electrons. The process of gaining electrons is called reduction, and a substance that gains electrons during a chemical reaction is said to be reduced. Oxidation and reduction occur together so that the total number of electrons lost by the substance that is oxidized will be equal to the number of electrons gained by the substance that is reduced. The number of moles of each reactant in the balanced chemical equation for an oxidation–reduction reaction reflects “electron balance” as well as “atom balance.”

The loss and gain of electrons by the reactants in Equation 1 will be more apparent if the overall reaction is broken down into two separate half-reactions. Equations 3 and 4 on the next page show the oxidation and reduction half-reactions, respectively, for the reaction of aluminum metal with copper(II) ions. In the oxidation half-reaction (Equation 3), each aluminum atom loses three electrons and is oxidized to an Al³⁺ ion. In the reduction half-reaction (Equation 4), each Cu²⁺ ion gains two electrons and is reduced to a copper atom. Notice that the total charge on both sides of a half-reaction must be the same (charge is conserved). The number of electrons involved in the overall oxidation–reduction reaction will be balanced if Equation 3 is multiplied by a factor of two (Equation 5a) and Equation 4 is multiplied by a factor of three (Equation 5b). When the resulting half-reactions are added together, the number of electrons lost by the aluminum atoms is equal to the number of electrons gained by the copper ions and the electrons “cancel out” of the overall equation. Equation 5c is the balanced, net ionic equation for the reaction of aluminum with copper(II) chloride—the chloride ions are “spectator ions” and are not shown.

Experiment Overview

The purpose of this experiment is to carry out a series of possible single replacement reactions of metals with solutions of metal cations in order to determine the activity series of the metals. The more active metal in each pair will be determined by observing which combinations of metals and metal cations undergo a chemical reaction. The metals will be ranked from most active to least active.

Materials

Prelab Questions

1. According to the Background information, magnesium is more active than aluminum. Predict which combination of metal and metal ion solutions will show evidence of a chemical reaction: (a) Mg(s) and Al₂(SO₄)₃(aq) or (b) Al(s) and MgSO₄(aq).
2. Write the balanced chemical equation for the single replacement reaction that is expected to occur (1a or 1b).
3. Which substance is being oxidized in this reaction? Which substance is being reduced?
4. Write the oxidation and reduction half-reactions for the overall reaction.

Safety Precautions

Silver nitrate is slightly toxic by ingestion and will stain skin and clothing. Copper(II) sulfate and iron(II) sulfate are toxic by ingestion. Metal pieces may have sharp edges—handle with care. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the lab.

Procedure

1. Place a 24-well reaction plate on top of a sheet of white paper as shown below. Each well is identified by a unique combination of a letter and a number, where the letter refers to a horizontal row and the number to a vertical column (see Figure 1).
   {12615_Procedure_Figure_1_Layout and numbering of a 24-well reaction plate}
2. Obtain or cut five 1-cm2 pieces of each metal to be tested—copper, iron, magnesium and zinc—and lay out the metals in labeled rows on a paper towel. Follow your instructor’s directions for polishing the metals, if necessary. Consult the data table frequently and carefully read each label. Notice that it is not necessary to test a metal with a solution of its own cation (e.g., copper is not tested in copper(II) sulfate).
3. Using a clean, labeled pipet for each solution, add 15 drops of the appropriate solution to each well, as follows:
   • Copper(II) sulfate solution to each well A1–D1
   • Iron(II) sulfate solution to each well A2–D2
   • Magnesium nitrate solution to each well A3–D3
   • Silver nitrate solution to each well A4–D4
   • Zinc sulfate solution to each well A5–D5
4. Place one piece of copper into each well A2–A5.
5. Place one piece of iron into well B1 and into each well B3–B5.
6. Place one piece of magnesium into each well C1, C2, C4 and C5.
7. Place one piece of zinc into each well D1–D4.
8. After all the metal pieces have been added, observe any evidence or “signs” of a chemical reaction in each well. Use a clean toothpick to submerge any metal pieces if necessary. Record all observations in the data table.
9. Some chemical reactions may be spontaneous but slow—continue to observe the reactions for 3–5 minutes, and record all observations in the data table.
10. Sometimes the only sign of a chemical reaction is a slight darkening or coating of the metal surface. Use clean forceps and a magnifier to observe any metals where the evidence of reaction is not obvious. It may be helpful to compare the color and appearance of a test metal with a clean, untreated piece of metal. Record all observations in the data table.
11. If no evidence of a reaction is observed in a well, write NR (no reaction). Remember, single replacement reactions will occur in one direction only. Many of the metal/metal ion combinations will show no reaction!
12. Do NOT dispose of metal pieces down the drain. Fold several paper towels together to form a thick mat. Gently shake the contents of the reaction plate onto the folded paper towels. Use forceps to remove any metal pieces sticking to the bottom or sides of the reaction plate. The folded paper towels with the metal pieces may be discarded in the trash.
13. Use cotton swabs to clean out any remaining residues in the reaction plate. Wash the reaction plate thoroughly under running water and rinse well with distilled water.

Student Worksheet PDF

12615_Student1.pdf