Materials Included In Kit

Additional Materials Required

Prelab Preparation

Part 1—Clean each of the metals (magnesium, copper and zinc) by rubbing each strip with the fine steel wool pad before cutting. Cut the magnesium metal ribbon into 1-cm strips. Cut the copper and zinc strips into small pieces, approximately 0.5 cm in width. Each metal strip should be cut into about 25–30 small pieces.

Part 2—Prepare the following solutions on the day of the lab. Prepare only as much as will be needed for the day. Instructions are provided for 30 mL each of bromine water and chlorine water—more than sufficient for a class of 30 students working in pairs. Prepare and dispense bromine and chlorine water in a working fume hood to avoid unnecessary exposure to vapors.

Bromine Water—Procedure for preparing 30 mL of bromine water: In a small labeled beaker, mix 10 mL of 1 M sodium bromide solution and 10 mL of 1 M hydrochloric acid solution. Add 10 mL of sodium hypochlorite solution and stir to mix the reactants. Place a pipet in the beaker for dispensing.

Chlorine Water—Procedure for preparing 30 mL of chlorine water: In a small labeled beaker, mix 4 mL of sodium hypochlorite solution with 20 mL of 1 M HCl solution. Dilute to 30 mL with distilled or deionized water. Place a pipet in the beaker for dispensing.

Safety Precautions

Conduct this lab in a fume hood or well-ventilated laboratory. Dispense the dilute bromine, chlorine, and iodine water solutions in an operating fume hood. Bromine water, chlorine water and hydrochloric acid are toxic by ingestion and inhalation. Bromine water is a strong oxidizing agent and a severe skin and eye irritant. Keep a supply of sodium thiosulfate solution on hand as a neutralizer when preparing or using bromine. Iodine solution is an eye-irritant and corrosive to the skin. Sodium hypochlorite solution is a corrosive liquid, causes skin burns and reacts with acids to evolve chlorine gas; moderately toxic by ingestion and inhalation. In this lab sodium hypochlorite is reacted with hydrochloric acid to generate small amounts of very dilute halogen solutions for use by the students. This step should only be performed by the teacher in the amounts indicated. Follow the directions carefully and work in an operating fume hood. Copper(II) sulfate solution is toxic by ingestion. Avoid contact of all chemicals with eyes and all body tissues. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please consult 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. Bromine water is readily neutralized by sodium thiosulfate according to Flinn Suggested Disposal Method #12a. Iodine solution can be disposed of according to Flinn Suggested Disposal Method #12a. Chlorine water can be disposed of by allowing it to stand in an open container inside an operating fume hood. The chlorine will slowly leave the water and the container. The remaining water (now effectively degassed) can be put down the drain. The metal ion solutions can be disposed of down the drain according to Flinn Suggested Disposal Method #26b. Metal pieces should be disposed of in the solid waste disposal.

Teacher Tips

• Enough materials are supplied in this kit for a class of 30 students working in pairs or for 15 groups of students. The lab can be reasonably completed in one 50-minute lab period.
• Consider having half of the class start on Part 1 and the other half start on Part 2 in order to decrease congestion at the reagent table.
• Warn students not to dump the solutions from the test tubes in Part 1 into the sink as the unreacted solid metal pieces may fall in and clog up the drain. Instead provide a liquid waste beaker for the metal ion solutions. You can then dispose of this down the drain with plenty of water (see Disposal section).
• Test tubes from Part 1 need to be rinsed out and reused for Part 2 unless additional test tubes are available in your lab.
• This is a qualitative rather than a quantitative lab activity. Amounts called for in the procedure are approximations and need not be exact. Amounts can be estimated by first filling a test tube with 10 drops of liquid and then filling the other tubes to the same height.

Sample Data

Data Table 1. Reactivity of the Metals

NR = No Reaction

Data Table 2. Halogens in Aqueous and Organic Solutions Data Table 3. Reactivity of the Halogens

Answers to Questions

Part 1. Reactivity of the Metals

1. Based on laboratory results, which metal is most active? How do you know this?

   The most active metal is magnesium since it replaced both of the other metals used in this experiment. (In other words, magnesium metal showed a chemical reaction with copper(II) ions and zinc ions.)

2. Which metal is least active? How do you know this?

   The least active metal is copper since it was replaced by both of the other metals used in this experiment. (In other words, copper metal showed no reaction with magnesium ions or zinc ions.)

3. Using lab data, rank the metals in order of reactivity, from most active to least active. What evidence do you have for this activity order?

   Magnesium, zinc, copper. Lab data showed that magnesium reacted with both metals, zinc only reacted with one and copper did not show any reaction at all.

4. Write a balanced equation for each reaction that occurred in Part 1. Write net ionic equations, omitting any spectator ions.

   Mg(s) + Cu²⁺(aq) → Cu(s) + Mg²⁺(aq)
   Mg(s) + Zn²⁺(aq) → Zn(s) + Mg²⁺(aq)
   Zn(s) + Cu²⁺(aq) → Cu(s) + Zn²⁺(aq)
   Zn(s) + Mg²⁺(aq) → NR
   Cu(s) + Mg²⁺(aq) → NR
   Cu(s) + Zn²⁺(aq) → NR

5. Look specifically at the reaction between magnesium metal and copper(II) ions.
   1. Write the half-reaction for magnesium metal. Is this an oxidation or a reduction reaction?

      Mg(s) → Mg²⁺(aq) + 2e^– Oxidation reaction

   2. Write the half-reaction for the copper(II) ions. Is this an oxidation or a reduction reaction?

      Cu²⁺(aq) + 2e^– → Cu(s)Reduction reaction

   3. Which material is oxidized and which is the oxidizing agent? Which material is reduced and which is the reducing agent?

      Magnesium is oxidized by the copper(II) ions (the oxidizing agent). Copper(II) ions are reduced by the magnesium metal (the reducing agent).

6. Which metal is most likely to retain electrons—magnesium, copper or zinc? Why?

   Copper, because copper did not react with magnesium or zinc ions, which shows it is least likely to give up its electrons.

7. Which metal has the greatest tendency to lose electrons? Why?

   Magnesium, because magnesium reacted with both copper and zinc ions, which shows that it has the greatest tendency to give up its electrons.

Part 2. Reactivity of the Halogens

8. Based on laboratory observations, how can each of the free halogens—Br₂, Cl₂ and I₂—be detected in the aqueous layer and in the organic (hexanes) layer?

   Elemental bromine is light orange/yellow in water and orange in hexanes. Elemental chlorine is colorless in both water and in hexanes. Iodine is dark orange in water and deep purple in hexanes.

9. Look at results from tubes 1 and 2. Did bromine react with the chloride ion solution? The iodide ion solution? What evidence do you have for any reaction that occurred?

   Bromine reacted with iodide ions but not with chloride ions. Evidence for reaction with iodide ions is the appearance of a purple color in the top organic layer (indicating elemental iodine) when bromine is mixed with iodide ions. No change occurs when bromine is mixed with chloride ions.

10. Look at results from tubes 3 and 4. Did chlorine react with the bromide ion solution? The iodide ion solution? What evidence do you have for any reaction that occurred?

    Chlorine reacted with both bromide and iodide ions. Evidence for reaction is the appearance of a purple color in the top organic layer (indicating elemental iodine) when chlorine is mixed with iodide ions, and appearance of an orange color in the top organic layer (indicating elemental bromine) when chlorine is mixed with bromide ions.

11. Look at results from tubes 5 and 6. Did iodine react with the bromide ion solution? The chloride ion solution? What evidence do you have for any reaction that occurred?

    Iodine did not react with either bromide or chloride ions as there was no change when mixed.

12. Write a balanced equation for each reaction that occurred in Part 2. Write net ionic equations, omitting any spectator ions.

    Br₂ + I^– → I₂ + Br^–
    Cl₂ + Br^– → Br₂ + Cl^–
    Cl₂ + I^– → I₂ + Cl^–
    Br₂ + Cl^– → NR
    I₂ + Br^– → NR
    I₂ + Cl^– → NR

13. Based on laboratory results, which halogen is the most active? Which is the least active? Provide evidence.

    Chlorine is the most active halogen while iodine is the least active. This is apparent by the observation that chlorine oxidized (reacted with) both bromide and iodide ions while iodine did not oxidize (react with) either bromide or chloride ions.

14. Using lab data, rank the halogens in order of their relative oxidizing ability, from highest to lowest. What evidence do you have for this activity order?

    Chlorine, bromine, iodine. Lab data showed that chlorine oxidized both halides, bromine only oxidized one, and iodine did not show any reaction at all.

15. Look specifically at the reaction between elemental bromine and iodide ions.
    1. Write the half-reaction for the elemental bromine. Is this an oxidation or a reduction reaction?

       Br₂ + 2e^– → 2Br^– Reduction reaction

    2. Write the half-reaction for the iodide ions. Is this an oxidation or a reduction reaction?

       2I^– → I₂ + 2e^– Oxidation reaction

    3. Which material is oxidized and which is the oxidizing agent? Which material is reduced and which is the reducing agent?

       Iodide ions are oxidized by the bromine (the oxidizing agent). Bromine is reduced by the iodide ions (the reducing agent).

Introduction

In this experiment, reactivity of the metals and of the halogens will be investigated. From laboratory observations, a relative metal activity series and relative halogen oxidizing ability will be developed.

Concepts

• Oxidation–reduction
• Metal reactivity
• Halogen reactivity

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 loses electrons and becomes more positive. 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 gains electrons and becomes more negative. 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 magnesium atoms to zinc ions. Magnesium donates electrons and its charge increases—it is oxidized. The zinc 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 zinc ion accepts electrons, so it is the oxidizing agent. In other words, the zinc ion oxidizes magnesium. Conversely, magnesium donates electrons, which makes it the reducing agent—it reduces the zinc ion.

Consider what is happening to the two reactants independently. Each magnesium atom loses two electrons to form a magnesium ion, as described by Equation 2. Magnesium donates electrons and becomes more positively charged. Equation 2 represents the oxidation half-reaction. For every oxidation, there must be a reduction. In this reaction, electrons are transferred to the zinc ions in solution to form zinc metal, as described by Equation 3. Zinc ions accept electrons and become reduced in charge. Equation 3 represents the reduction half-reaction. Metal Reactivity
A common type of redox reaction is a single replacement reaction. Single replacement reactions involve the replacement of one element in a compound with another element. The general form for single replacement reactions is shown in Equation 4. Metals are commonly involved in single replacement reactions. Some metals can replace other metals in their compounds, while some metals cannot. In Equation 4, if “A” and “B” are metals, “A” replaces “B” in its compound “BC.” The reaction is not reversible, so “B” cannot replace “A” in the compound “AC.” The ability to replace another metal determines a metal’s reactivity—the better the ability to replace another metal, the more reactive a metal is. The activity series of metals is a list that places the metals in order of reactivity. The metals at the top are the most reactive and can therefore replace most other metals. Reactivity decreases as you move down the list, with those at the bottom of the list capable of replacing only a few other metals. A metal can replace another metal if it appears above that metal in the activity series. In Equation 4, “A” must be the more reactive metal (higher up on the activity series and capable of replacing “B”) while “B” is the less active metal (lower on the activity series and not capable of replacing “A”). The activity series for metals is determined by experiments in which pairs of metals are compared for reactivity.

Halogen Reactivity
Halogens are also commonly involved in single replacement reactions. Some halogens have a greater oxidizing ability and are therefore more reactive. In Equation 5, if X and Y are halogens and M is a metal cation, then X₂ in the uncombined (free element) form will oxidize Y^– to Y₂. After the reaction has occurred, the X is now combined with the metal M, and Y exists as a free element. The reaction is usually described with the words, X has replaced Y. The reaction is not reversible; in other words, the reaction written in Equation 6 will not take place. The ability to oxidize or replace another halogen determines a halogen’s reactivity—the better the ability to oxidize another halogen, the more reactive is the halogen. A reactivity series can be devised which lists the halogens in order from the most reactive to the least reactive halogen. The activity series for halogens is determined by experiments in which pairs of elemental halogens and halide ions are compared for reactivity.

Materials

Safety Precautions

Conduct this lab in a fume hood or well-ventilated laboratory. Bromine water and chlorine water are toxic by ingestion and inhalation. Bromine water is a strong oxidizing agent and a severe skin and eye irritant. Iodine solution is an eye-irritant and somewhat corrosive to the skin. Copper(II) sulfate solution is toxic by ingestion. Avoid contact of all chemicals with eyes and skin. Do not allow sodium hypochlorite (bleach) to come in contact with acids—toxic chlorine gas may be generated. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part 1. Reactivity of the Metals: An activity series for three metals will be determined by combining pairs of metals with metal ions and comparing their reactivity.

1. Place six clean test tubes in a test tube rack. Label the tubes 1–6.
2. Obtain two small pieces of each of the three metals to be tested—magnesium, copper and zinc.
3. Place one piece of magnesium into tubes 1 and 2.
4. Place one piece of copper into tubes 3 and 4.
5. Place one piece of zinc into tubes 5 and 6.
6. Add 10 drops of 0.1 M CuSO₄ solution to the Mg in tube 1. Observe any reaction that occurs—bubbles, fizzing, color changes or any change that occurs on the surface of the metal. Record detailed observations in Data Table 1. If no reaction is apparent, write no reaction (NR) in the data table.
7. Continue to add 10 drops of the appropriate 0.1 M solution to the metals in the tubes as indicated in the chart below. Make and record detailed observations as you add each solution.
   {13363_Procedure_Table_1}
8. Dispose of the solutions by decanting the solution from each tube down the drain with plenty of water. Empty the leftover waste metal in a solid waste container provided by your teacher. Rinse the tubes with water for use in Part 2.

Part 2. Reactivity of the Halogens: An activity series for three halogens will be determined by combining elemental halogens with halide ions and comparing their reactivity. The halogens—bromine, chlorine and iodine—will be identified by their solubility and unique colors in both water and an organic solvent, hexane.

9. Place six clean test tubes in a test tube rack. Label the tubes 1–6.
10. Use a pipet to add 10 drops of bromine water (Br₂) to tubes 1 and 2.
11. Use a different pipet to add 10 drops of chlorine water (Cl₂) to tubes 3 and 4.
12. Use a different pipet to add 10 drops of iodine solution (I₂) to tubes 5 and 6.
13. Record the color of each aqueous free halogen solution in Data Table 2.
14. Add approximately 10 drops of hexanes (organic solvent) to each tube (1–6).
15. Place a stopper on each tube and carefully invert each tube several times to mix the liquids. Notice that there are now two layers—the top organic layer (hexanes) and the bottom aqueous layer (water).
16. For each halogen, record the color of the upper organic layer and the lower aqueous layer in Data Table 2.
17. Add 10 drops of 0.1 M NaCl solution to the Br₂ in tube 1. Replace the stopper on the tube and carefully invert the tube several times to mix the liquids. Observe any reaction that occurs—look for changes in color to either layer. Record the color changes in the Data Table 3. If no change is apparent, write no reaction (NR) in the data table.
18. Continue to add 10 drops of the appropriate 0.1 M solution to the tubes as indicated in the following chart. Make and record detailed observations as you add each solution.
    {13363_Procedure_Table_2}
19. Dispose of the solutions by pouring the contents of the test tubes into an appropriate liquid waste container supplied by your teacher. Rinse the tubes with water for later use.

Student Worksheet PDF

13363_Student1.pdf