Materials Included In Kit

Additional Materials Required

Prelab Preparation

1. Prepare 1 M sodium bicarbonate solution by dissolving 2.1 g of solid sodium bicarbonate in 25 mL of distilled or deionized water.
2. Cut the magnesium ribbon into 3-cm pieces for student use. Each student group will use four pieces in Part I.

Safety Precautions

Hydrochloric acid solution is toxic by ingestion and inhalation. It is corrosive to skin and eyes. Avoid body tissue contact. Ammonium hydroxide solution is very irritating to the skin and eyes. It is toxic by ingestion and inhalation. Avoid body tissue contact. Use in a fume hood. Sodium hydroxide solution is corrosive to skin and eyes. Avoid body tissue contact. Magnesium is a flammable solid and burns with an intense flame. Silver nitrate solution is mildly toxic by ingestion and will discolor skin and clothing. Avoid skin contact. Copper(II) sulfate solution is a skin irritant. Please review current Safety Data Sheets for additional safety, handling and disposal information. Wear chemical splash goggles, chemical-resistant gloves and apron. Have students wash hands thoroughly with soap and water before leaving the laboratory.

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. Sodium hydroxide and ammonium hydroxide solutions may be neutralized according to Flinn Suggested Disposal Method #10. Hydrochloric acid solution may be neutralized according to Flinn Suggested Disposal Method #24b. Sodium bicarbonate may be disposed of according to Flinn Suggested Disposal Method #26a. Silver ion may be precipitated according to Flinn Suggested Disposal Method #11. Potassium iodide and the copper sulfate solutions may be disposed of according to Flinn Suggested Disposal Method #26b.

Lab Hints

• The entire lab will take two 50-minute lab periods to complete. The end of Part II is a convenient stopping point.

Part I

• In step 4, the bubbling of the magnesium is somewhat unexpected. The copper(II) sulfate solution is acidic. There are two competing single-replacement reactions taking place.

Part II

• Step 7 shows a light blue precipitation from the OH^– which is quickly overpowered by the formation of the complex ion as more ammonium hydroxide is added
• Step 9 seems to produce a mixture of Cu(OH)₂ precipitate and CuCO₃ precipitate.
• Step 10 is somewhat unexpected. A precipitate is not expected from the formation of copper(II) iodide or potassium sulfate since both are soluble. As one examines the reagents looking for possible reaction, the oxidation–reduction possibility is found. Cu²⁺ can reduce to Cu⁺ or Cu depending on the reducing agent. I^– is a medium-strength reducing agent. Thus, it is expected that Cu²⁺ will reduce to Cu⁺ but the presence of I^– (even as some has oxidized to I₂) will cause a precipitate of CuI.

Part III

• In step 14, the formation of water cannot be observed. Some heat can be detected.

Part IV

• Step 15 is difficult to predict—the final product could be sodium carbonate or sodium oxide, and some students might even predict sodium hydroxide. Because a bicarbonate is involved, carbon dioxide is expected to be given off. Sodium carbonate as it forms will not decompose easily in a normal burner flame since its melting point is 851 °C.

Teacher Tips

• A discussion of the nomenclature of chemical compounds is useful prior to the assigning of a prelab. Expect that most students will have difficulty in predicting the reactions. Other methods of analysis can be discussed. Single replacement and double replacement reactions can be useful but it is necessary to emphasize that net ionic equations are to be written.
• For precipitation reactions, students should refer to a solubility table.

Answers to Prelab Questions

Read the lab Procedure before beginning. All chemicals used have been written in word form so that the formulas of the reagents must be determined. In all cases, a chemical reaction will occur. Write the formulas for all reactants, followed by their physical states—(g), (l), (s) or (aq)—for each step in the procedure. Predict the type of reaction (reaction class) for each step, and write an equation for each predicted reaction. Do not balance the reaction equations. (Reacting ionic species should be written as ions and reacting molecular species should be written as molecules.)

Sample Data

References

This lab was developed by DeWayne Lieneman, chemistry instructor, retired, Indian Head Park, Illinois.

Introduction

What evidence may be used to suggest that two reactants have generated a new product in a chemical reaction? The purpose of this lab is to predict the products when pairs of reactants are combined and to use observational data to confirm those predictions.

Concepts

• Reactants and products
• Chemical equations
• Net ionic equations
• Precipitation reactions
• Acid–base reactions
• Oxidation–reduction reactions
• Complex ion formation reactions

Background

The actual chemical species in a solution may not be the same as the written formula. Chemicals involved in reactions may be written in molecular or ionic form. Consider HCl. This could be hydrogen chloride, HCl(g), where there are molecules of gaseous HCl. If HCl is in an aqueous solution, HCl(aq), the molecule has dissociated into H⁺(aq) and Cl^–(aq) ions. Thus a chemical reacting with hydrochloric acid solution is not really reacting with HCl molecules, but rather with H⁺(aq) ions or Cl^–(aq) ions.

The essence of any chemical reaction is summarized in the form of a chemical equation. Consider the reaction represented by Equation 1, the burning of natural gas (methane) in a laboratory burner.

The reactants—or more specifically, their formulas—are written on the left side of the equation, the products on the right side of the equation. An arrow represents the direction of the reaction and is read as “yields” or “produces.” Other symbols are used to describe the physical state of the reactants and products (see Table 1). Molecular equations are often used to balance reactions or to determine quantities of reactants and products. Two examples of molecular equations are listed. In Equation 2, the reaction involves the molecules of gaseous nitrogen, NO(g), and gaseous chlorine, Cl₂(g). In Equation 3, the actual reaction is between the Pb²⁺(aq) ions and the Cl^–(aq) ions. The Na⁺(aq) ions and the NO₃ ^–(aq) ions are not involved in the reaction and are considered spectator ions. When Equation 3 is written with the species that occur in solution, it becomes Since the numbers of sodium ion, Na⁺(aq), and nitrate ions, NO₃ ^–(aq), are the same on both sides of the equation, the equation can be reduced to the net ionic equation. This lists only the ions and molecules involved in the chemical reaction. Prediction of reaction products can be difficult. To simplify the process, look for types of reactions in a particular order. In this lab, four categories of chemical reactions will be examined: oxidation–reduction, acid–base, decomposition and complex-ion formation.

Reactions involving the processes of oxidation and reduction are quite common. Oxidation–reduction reactions always occur together. Oxidation occurs when one reactant loses electrons and reduction occurs when another reactant gains those electrons.

To determine if a reaction is a redox reaction, first examine the reactants. Are any of the reactants elements? If elements are involved in reactions, they undergo electron transfer to form compounds. Some examples are chlorine gas, Cl₂(g); oxygen, O₂(g); iron, Fe(s); and sodium metal, Na(s). Are there metals present in ionic form in solution, either as a separate ion or as part of another ion? Metals in solution frequently undergo an oxidation state change in a chemical reaction. Some examples are iron(II) ion, Fe²⁺(aq); chromium(III) ion, Cr³⁺(aq); dichromate ion, Cr₂O₇^2–(aq); and permanganate ion, MnO₄ ^–(aq).

Now look at the products. Has a metal either been produced or consumed as a result of the reaction? This is always a sign of the occurrence of oxidation–reduction. Have the reactants changed oxidation states? Examine, if possible, the oxidation states of the atoms in the reactants and products. Have any changed? If so, the reaction involves oxidation–reduction. Examples are listed in Table 2. If the reaction doesn’t appear to be a redox reaction, next determine it may be an acid–base reaction. In an acid–base reaction, there must be an acid, with the general formula of HA, as a reactant, along with a base, with the general formula of MOH. The ionic equation for a general acid–base reaction is The net ionic equation is Examples of acid–base reactions are listed in Table 2.

If the reaction is neither redox or acid–base, two additional categories to consider are decomposition or precipitation reactions. If only one compound occurs as a reactant, the reaction is probably a decomposition of that compound. The products will be two or more compounds. If the reaction occurs in solution and involves two salts, then the reaction if probably a precipitation. The products of this reaction are a precipitate and a soluble salt. Examples are listed in Table 2.

Finally, look to see if the reactants could produce a complex ion in solution. This reaction usually has a metal ion present as one reactant (i.e., Fe³⁺(aq), and a complexing anion as the other reactant). Three common complexing ions are cyanide ion CN^–(aq), hydroxide ion OH^–(aq), and thiocyanate ion SCN^–(aq). Neutral molecules, such as NH₃ (ammonia), which can donate lone pairs of electrons to the metal ion, may also serve as complexing agents. The product of a complex-ion reaction is a single ion in solution that is a combination of the two reactant ions. The oxidation state of the metal does not change and the solution often undergoes color change. Examples are listed in Table 2.

Experiment Overview

The purpose of this experiment is to determine formulas of the reactants of possible chemical reactions and predict their products. The reactions are then run and evidence of reaction is observed. Those reactions are then classified as oxidation–reduction, acid–base, decomposition, precipitation or complex-ion reactions. The net ionic equations for these reactions will also be determined.

Materials

Safety Precautions

Hydrochloric acid solution is toxic by ingestion and inhalation. It is corrosive to skin and eyes. Avoid body tissue contact. Ammonium hydroxide solution is very irritating to the skin and eyes. It is toxic by ingestion and inhalation. Avoid body tissue contact and use in a fume hood. Sodium hydroxide solution is corrosive to skin and eyes. Avoid body tissue contact. Magnesium is a flammable solid and burns with an intense flame. Silver nitrate solution is mildly toxic by ingestion and will stain skin and clothing. Avoid skin contact. Copper(II) sulfate solution is a skin irritant. Wear chemical splash goggles, chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part I. Magnesium

1. Obtain a 3-cm piece of magnesium ribbon. Hold it with tongs in a burner flame. Caution: Do not look directly at the burning magnesium. Record all observations in the Data Table.
2. (a) Place 50 mL of water in a 150-mL beaker and heat to boiling. (b) Remove the water from heat. (c) Clean a 3-cm strip of magnesium ribbon. (d) Using tongs, drop it into the hot water. (e) Add 1–2 drops phenolphthalein. Observe again at the end of the period. Record all observations in the Data Table.
3. Obtain a 13 x 100 mm test tube. With a Beral-type pipet, add about 5 mL of the hydrochloric acid solution. Add a 3-cm strip of magnesium ribbon and observe. Record all observations in the Data Table.
4. Take a 3-cm strip of magnesium ribbon and clean it. Add it to a 13 x 100 mm test tube containing about 2 mL of the copper sulfate solution. Record all observations in the Data Table.

Part II. Copper Sulfate 

5. Using a Beral-type pipet, add 2 mL of the sodium phosphate solution to a 13 x 100 mm test tube. Add 2 mL of the copper sulfate solution to the test tube. Record all observations in the Data Table.
6. In a 20 x 150 mm test tube, mix 5 mL of the copper sulfate solution with 5 mL of the sodium hydroxide solution. Record all observations in the Data Table. Pour the contents into an evaporating dish and save for step 8.
7. In a 20 x 150 mm test tube, mix 2 mL of the copper sulfate solution with 5 mL of the ammonium hydroxide solution. Record all observations in the Data Table.
8. Pour about half of the mixture from the evaporating dish into a 20 x 150 mm test tube for use in step 9. Now carefully heat the evaporating dish to dryness. Record all observations in the Data Table.
9. Add 3 mL of the hydrochloric acid solution to the test tube saved from step 8. Record all observations in the Data Table.
10. Using a Beral-type pipet, add 1 mL of the copper sulfate solution and 1 mL of potassium iodide solution to a 13 x 100 mm test tube. Record all observations in the Data Table.

Part III. Hydrochloric Acid 

11. Place about 0.2 g of solid sodium bicarbonate in a 13 x 100 mm test tube. Using a Beral-type pipet, add 1 mL of hydrochloric acid solution. Record all observations in the Data Table.
12. In a 13 x 100 mm test tube, mix about 1 mL of the sodium bicarbonate solution and 1 mL of the hydrochloric acid solution. Record all observations in the Data Table.
13. In a 13 x 100 mm test tube, mix about 1 mL of the silver nitrate solution and 1 mL of the hydrochloric acid solution. Record all observations in the Data Table.
14. In a 13 x 100 mm test tube, mix about 1 mL of the sodium hydroxide solution and 1 mL of the hydrochloric acid solution. Record all observations in the Data Table.

Part IV. Miscellaneous 
Caution: Silver nitrate reacts with light and will stain skin. Avoid contact with skin and clothing.

15. In a 20 x 150 mm Pyrex^® test tube, heat about 2 g of solid sodium bicarbonate. Record all observations in the Data Table.
16. In a 13 x 100 mm test tube, mix about 1 mL of the potassium iodide solution and 1 mL of the silver nitrate solution. Record all observations in the Data Table.
17. In a 13 x 100 mm test tube, mix about 1 mL of the sodium hydroxide solution and 1 mL of the silver nitrate solution. Record all observations in the Data Table.
18. Add 2 mL of ammonium hydroxide solution to the mixture in the test tube from step 17.

Student Worksheet PDF

13017_Student1.pdf