Materials Included In Kit

Additional Materials Required

Prelab Preparation

Sodium phosphate solution, 0.5 M: Dissolve 9.5 g of sodium phosphate dodecahydrate (Na₃PO₄•12H₂O) in 50 mL of distilled or deionized water.

Safety Precautions

Ethyl alcohol is a flammable solvent and a dangerous fire risk. Keep away from heat, sparks and open flames. Solvent bottles should be kept capped at all times and must be removed from the work area when using the laboratory burner (Reaction 8). Addition of denaturant makes ethyl alcohol poisonous. It is harmful if swallowed and may cause skin, eye and respiratory irritation. Hydrochloric acid and sodium hydroxide solutions may cause severe skin burns and eye irritation. Keep sodium carbonate and citric acid on hand to clean up acid and base spills, respectively. Phenolphthalein is an alcohol-based solution and a flammable liquid; it is harmful if swallowed and may cause cancer or reproductive toxicity. Do not use this solution if you are pregnant. Ammonium carbonate causes skin and eye irritation and may cause respiratory irritation. Magnesium metal is a flammable solid and zinc metal may contain flammable dust. Copper(II) chloride solution is toxic by ingestion and may be harmful if swallowed. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or 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

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. Excess sodium hydroxide solution may be neutralized according to Flinn Suggested Disposal Method #10. Excess hydrochloric acid may be neutralized according to Flinn Suggested Disposal Method #24b. The mixture resulting from Reaction 5 contains solid copper metal and will clog the drains if discarded in the sink. Filter this mixture and discard the solid in the trash according to Flinn Suggested Disposal Method #26a. All other reaction mixtures may be rinsed down the drain with excess water according to Flinn Suggested Disposal Method #26b.

Lab Hints

• The goal of this survey experiment is for students to develop “chemical literacy” skills by observing as many chemical reactions as is practical. The laboratory work for the experiment can easily be completed within a typical 2-hour lab period. Student preparation is an essential element of lab safety—students should complete the Prelaboratory Assignment before class.
• Do not do the combustion of ethyl alcohol reaction (8) unless a designated work area can be set up in an isolated area of the lab, away from flammable solvents. For a safer alternative example of combustion, simply set up a 250-mL beaker with cold water on a ring stand above a laboratory burner. Light the burner and immediately observe water condensing on the outside of the beaker.
• Additional reactions may be assigned, if desired, to give students more experience recognizing and interpreting chemical reactions. Examples include burning sulfur to produce SO₂, single-replacement reaction of iron metal with copper(II) chloride, complex-forming combination reaction of copper(II) ions with ammonia, and the catalyzed decomposition of hydrogen peroxide.
• To ease congestion and improve efficiency carrying out this experiment, set up eight stations around the lab, one station per reaction. Students may rotate through the stations in any order. “Down time” between stations may be used to record observations.
• For safety reasons, we recommend using porcelain evaporating dishes to test the flammability of ethyl alcohol. If glass evaporating dishes or watch glasses are substituted, make sure all glassware is made from heat-resistant borosilicate (e.g., Pyrex^®) glass. Check the glassware for cracks and chips before use.
• This experiment provides experience in recognizing chemical reactions, writing chemical equations, and classifying chemical reactions. The best time to do this experiment is prior to lecture instruction on the topic. Students who get the theory before the practice tend to think of chemical reactions as formulas on a piece of paper. As discussed in the Background section, chemical reactions are quite showy events, with lots of interesting observations. In most cases, the observations themselves will enable students to identify or predict the products from the known structures of the reactants.
• Different textbooks may classify reactions differently. Some authors treat all single replacement reactions within a much larger group of oxidation–reduction reactions. Other authors refer to double replacement reactions as metathesis or double displacement reactions. Finally, a few authors use neither designation and classify double replacement reactions as either precipitation or acid–base reactions.
• Students may not be familiar with the properties of the common gases that are used to identify the gaseous products in reactions 2, 3 and 4. See the “Properties and Identification of Common Gases” in the Further Extensions section for a convenient summary and student handout.
• Demonstrations are a classic—and very popular—way to reinforce the exciting variety of chemical reactions. Reaction of calcium oxide with water may be used to show a highly exothermic combination reaction while microscale electrolysis of water illustrates a decomposition reaction. The single replacement reaction of aluminum with copper(II) chloride is a wonderful activity to use to develop students’ observation skills.

Further Extensions

Supplementary Information

Properties and Identification of Common Gases
Many common gases can be identified by testing their combustion properties (the burning splint test) and/or their acid–base properties (the litmus test).

• Carbon dioxide gas (CO₂) is more dense than air and does not support combustion. If a burning wood splint is inserted into a container of carbon dioxide, the flame will be smothered or extinguished due to the lack of oxygen. Carbon dioxide gas dissolves in water to give an acidic solution, and can be detected using a saturated limewater solution, Ca(OH)₂(aq). Carbon dioxide combines with calcium hydroxide to form calcium carbonate, which precipitates from solution.
  {14029_Extensions_Reaction_5}
• Oxygen gas (O₂) supports combustion and is necessary for combustion reactions to occur. If a glowing (not burning) splint is placed in an oxygen atmosphere, it will reignite and burn brightly.
• Hydrogen gas (H₂) is less dense than air and is flammable. It forms an explosive mixture with air. When a burning splint is exposed to hydrogen gas, a loud “pop” is usually heard as the hydrogen explodes and the flame is extinguished. The product of the combustion reaction of hydrogen is water. Hydrogen is essentially insoluble in water.

  2H₂(g) + O₂(g) → 2H₂O(l)

• Ammonia gas (NH₃) has a characteristic odor and dissolves readily in water to give a basic solution. Ammonia can be detected by placing a piece of moist litmus paper in the stream of ammonia gas released from a reaction mixture. The litmus paper turns blue (basic) due to the formation of hydroxide ions, as shown.
  {14029_Extensions_Reaction_6}
• Sulfur dioxide (SO₂) and sulfur trioxide (SO₃) are colorless gases with stinging odors. They dissolve in water to form acidic solutions.
  {14029_Extensions_Reaction_7}

Answers to Prelab Questions

1. Which reactants used in this experiment are flammable? Discuss the safety precautions that are necessary when working with flammable chemicals.

   Ethyl alcohol is a flammable solvent. Phenolphthalein solution contains alcohol and is also flammable. Keep away from flames and other sources of ignition. Cap all solvent bottles and remove them, along with all other flammable materials, from the designated work area when doing the burning alcohol test.

2. Summarize the following description of a chemical reaction in the form of a balanced chemical equation.

   “When solid sodium bicarbonate is heated in a test tube, an invisible gas, carbon dioxide, is released into the surrounding air. Water condenses at the mouth of the test tube and a white solid residue, sodium carbonate, remains behind in the bottom of the test tube.”
   2NaHCO₃(s) → Na₂CO₃(s) + H₂O(l) + CO₂(g)

3. Common observations of a chemical reaction are described in the Background section. For each observation, name a common or everyday occurrence that must involve a chemical reaction. Example: When a candle burns, it gives off light and heat. The production of light and heat is evidence for a chemical reaction.
   {14029_PreLabAnswers_Table_3}
4. The relative reactivity, or activity, of different metals can be determined by studying their single-replacement reactions with other metal compounds. Rank the following metals, copper, zinc and silver, from most active to least active based on the reactions summarized. Explain your reasoning.

   Zn(s) + CuCl₂(aq) → Cu(s) + ZnCl₂(aq)
   Cu(s) + ZnCl₂(aq) → NR
   Cu(s) + 2AgNO₃(aq) → 2Ag(s) + Cu(NO₃)₂(aq)
   Activity series: Zn > Cu > Ag
   Zinc is more active than copper because it reacted with copper(II) chloride, whereas copper metal did not react with zinc chloride. Silver is the least active because copper metal reacted with silver nitrate. (We assume that the reverse reaction would not take place—it doesn’t.)

Sample Data

Answers to Questions

1. Write a balanced chemical equation for each reaction 1–8. Classify each reaction using the information provided in the Background section (see Table 2).
   {14029_Answers_Table_5}
2. Classifying chemical reactions helps chemists to predict the possible products that will form when two or more substances are mixed. Complete and balance the following equations by predicting the products of each chemical reaction.
   1. Double replacement: 2NaOH(aq) + CuSO₄(aq) → Cu(OH)₂(s) + Na₂SO₄(aq)
   2. Combination: 4Al(s) + 3O₂(g) → 2Al₂O₃(s)
   3. Combustion: C₆H₁₂O₆(s) + 6O₂(g) → 6CO₂(g) + 6H₂O(l)
   4. Decomposition: CaCO₃(s) → CaO(s) + CO₂(g)
   5. Single replacement: 2Fe(s) + 3Pb(NO₃)₂(aq) → 3Pb(s) + 2Fe(NO₃)₃(aq)
3. Write a balanced chemical equation for each reaction and classify the reaction.
   1. Copper metal heated with oxygen gives solid copper(II) oxide.
      {14029_Answers_Reaction_1}
   2. Mixing ammonium nitrate and sodium hydroxide solutions gives aqueous sodium nitrate, ammonia gas and water.

      NH₄NO₃(aq) + NaOH(aq) → NaNO₃(aq) + NH₃(g) + H₂O(l) Double replacement

   3. Mercury(II) nitrate solution reacts with potassium iodide solution to give a mercury(II) iodide precipitate and potassium nitrate solution.

      Hg(NO₃)₂(aq) + 2KI(aq) → HgI₂(s) + 2KNO₃(aq) Double replacement

   4. Aluminum metal and sulfuric acid yield aqueous aluminum sulfate and hydrogen gas.

      2Al(s) + 3H₂SO₄(aq) → Al₂(SO₄)₃(aq) + 3H₂(g) Single replacement

   5. Acetic acid and lithium hydroxide solution produce water and aqueous lithium acetate.

      HC₂H₃O₂(aq) + LiOH(aq) → H₂O(l) + LiC₂H₃O₂(aq) Double replacement

   6. Sulfur dioxide gas reacts with oxygen on a platinum catalyst surface to produce sulfur trioxide gas.
      {14029_Answers_Reaction_2}
   7. Sodium metal reacts with water to give sodium hydroxide solution and hydrogen gas.

      2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) Single replacement

   8. Heating solid nickel chloride dihydrate yields solid nickel chloride and water vapor.
      {14029_Answers_Reaction_3}
   9. Heating solid potassium chlorate in the presence of manganese dioxide catalyst produces potassium chloride and oxygen gas.
      {14029_Answers_Reaction_4}

Introduction

The power of chemical reactions to transform our lives is visible all around us—in our homes, in our cars, even in our bodies. Chemists try to make sense of the great variety of chemical reactions the same way that biologists organize their knowledge of life, by sorting reactions into groups and classifying them. Classifying chemical reactions allows us to predict what chemical reactions will occur when different substances are mixed.

Concepts

• Chemical reactions
• Combination vs. decomposition
• Single vs. double replacement
• Combustion reactions

Background

A chemical reaction is defined as any process in which one or more substances are converted into new substances with different properties. Chemical reactions change the identity of the reacting substance(s) and produce new substances. Observing the properties of the reactants and products is therefore a key step in identifying chemical reactions. Some of the observations that may be associated with a chemical reaction include: (1) release of a gas; (2) formation of a precipitate; (3) color changes; (4) temperature changes; (5) emission or absorption of light. As these observations suggest, chemical reactions can be dynamic and exciting events. The essence of any chemical reaction—reactants being transformed into products—is summarized in the form of a chemical equation. Consider the reaction represented by Equation 1, the burning of natural gas (methane, CH₄) in a laboratory burner.

The reactants—or, more specifically, their formulas—are written on the left side of the equation, and 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 may be used to describe the physical state of the reactants and products and to describe the reaction conditions (see Table 1). Chemical reactions may be classified by considering the number of reactants and products in the reaction, the physical or chemical nature of the reactants and products and the rearrangement of atoms in the conversion of the reactants into products (see Table 2).

Experiment Overview

The purpose of this experiment is to observe a variety of chemical reactions and to identify patterns in the conversion of reactants into products. The properties of the reactions will be analyzed to classify the chemical reactions into different groups.

Materials

Prelab Questions

1. Which reactants used in this experiment are flammable? Discuss the safety precautions that are necessary when working with flammable chemicals.
2. Summarize the following description of a chemical reaction in the form of a balanced chemical equation.

   “When solid sodium bicarbonate is heated in a test tube, an invisible gas, carbon dioxide, is released into the surrounding air. Water condenses at the mouth of the test tube and a white solid residue, sodium carbonate, remains behind in the bottom of the test tube.”

3. Common observations of a chemical reaction are described in the Introduction section. For each observation, name a common or everyday occurrence that must involve a chemical reaction. Example: When a candle burns, it gives off light and heat. The production of light and heat is evidence for a chemical reaction.
4. The relative reactivity, or activity, of different metals can be determined by studying their single-replacement reactions with other metal compounds. Rank the following metals, copper, zinc and silver, from most active to least active based on the reactions summarized below. Explain your reasoning.

   Zn(s) + CuCl₂(aq) → Cu(s) + ZnCl₂(aq)
   Cu(s) + ZnCl₂(aq) → NR
   Cu(s) + 2AgNO₃(aq) → 2Ag(s) + Cu(NO₃)(aq)

Safety Precautions

Ethyl alcohol is a flammable solvent and a dangerous fire risk. Keep away from heat, sparks and open flames. Solvent bottles should be kept capped at all times and must be removed from the work area when using the laboratory burner (Reaction 8). Addition of denaturant makes ethyl alcohol poisonous. It is harmful if swallowed and may cause skin, eye and respiratory irritation. Hydrochloric acid and sodium hydroxide solutions may cause severe skin burns and eye irritation. Notify the instructor and clean up all spills immediately. Phenolphthalein is an alcohol-based solution and a flammable liquid; it is harmful if swallowed and may cause cancer or reproductive toxicity. Do not use this solution if you are pregnant. Ammonium carbonate causes skin and eye irritation and may cause respiratory irritation. Magnesium metal is a flammable solid and zinc metal may contain flammable dust. Copper(II) chloride solution is toxic by ingestion and may be harmful if swallowed. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the lab.

Procedure

For each reaction, record the color and appearance of the reactant(s), the evidence for a chemical reaction and the properties of the product(s).

Reaction 1

1. Obtain a 3–4 cm strip of magnesium metal ribbon. Hold the piece of magnesium with forceps or crucible tongs and heat the metal in a laboratory burner flame. Caution: Do not look directly at the burning magnesium—ultraviolet light is produced that may damage your eyes.
2. When the magnesium ignites, remove it from the flame and hold it over an evaporating dish or a Pyrex^® watch glass until the metal has burned completely. Let the product fall into the evaporating dish.
3. Turn off the laboratory burner and observe the properties of the product in the evaporating dish.

Reaction 2

4. Using a Beral-type pipet, add about 2 mL (40 drops) of 1 M hydrochloric acid solution to a small test tube.
5. Obtain a 2–3 cm strip of magnesium metal ribbon and coil it loosely into a small ball. Add the magnesium metal to the acid in the test tube.
6. Carefully feel the sides of the test tube and observe the resulting chemical reaction for about 30 seconds.
7. While the reaction is still occuring, light a wood splint and quickly place the burning splint in the mouth of the test tube. Do not put the burning splint into the acid solution.

Reaction 3

8. Obtain a clean, dry test tube and place a small amount (about the size of a jelly bean) of ammonium carbonate into the test tube.
9. Use a test tube clamp to hold the test tube and gently heat the tube in a laboratory burner flame for about 30 seconds.
10. Remove the test tube from the flame and place a piece of moistened litmus paper in the mouth of the test tube. Identify any odor that is readily apparent by wafting the fumes toward your nose. Caution: Do NOT sniff the test tube!
11. Test for the formation of a gas: Light a wood splint and insert the burning splint halfway down into the test tube.

Reaction 4

12. Place a small amount (about the size of a jelly bean) of calcium carbonate into a clean, dry test tube.
13. Using a Beral-type pipet, add about 1 mL (20 drops) of 1 M hydrochloric acid to the test tube. Feel the sides of the test tube and observe the reaction for 30 seconds.
14. Light a wood splint and insert the burning splint about halfway down into the test tube. Do not allow the burning splint to contact the reaction mixture.

Reaction 5

15. Using a Beral-type pipet, add about 2 mL (40 drops) of 0.5 M copper(II) chloride solution into a small test tube.
16. Add 1–2 pieces of mossy zinc or one piece of zinc shot to the test tube and observe the resulting chemical reaction.

Reaction 6

17. Using a Beral-type pipet, add about 2 mL (40 drops) of 0.5 M copper(II) chloride solution into a small test tube.
18. Using a fresh pipet, add about 25 drops of 0.5 M sodium phosphate solution to the test tube.

Reaction 7

19. Using a Beral-type pipet, add 20 drops of 1 M sodium hydroxide solution to a small test tube.
20. Add one drop of phenolphthalein indicator to the test tube and mix the solution by gently swirling the tube. Hint: Phenolphthalein is called an “acid–base” indicator.
21. Using a clean Beral-type pipet, add 1 M hydrochloric acid solution one drop at a time to the test tube. Count the number of drops of acid required for a permanent color change to be observed.

Reaction 8

22. Working in the hood or a designated work area, add about 1 mL (20 drops) of ethyl alcohol to a clean porcelain evaporating dish. Place the evaporating dish on a heat-resistant pad.
23. Cap the alcohol bottle and remove it from the work area.
24. Fill a test tube about one-third full with cold tap water for use in step 27.
25. Light a butane safety lighter and bring the flame close to the alcohol in the evaporating dish.
26. Extinguish the safety lighter as soon as the alcohol ignites.
27. Place the test tube containing cold water in a test tube clamp and hold the test tube above the burning alcohol. Observe the outside of the test tube for evidence of product formation.
28. Allow the alcohol to burn until it is completely consumed. Caution: Do not touch the hot evaporating dish.

Student Worksheet PDF

14029_Student1.pdf