Classifying Chemical Reactions

Introduction

Use this series of colorful and dramatic demonstrations to teach and then assess student understanding of five basic types of chemical reactions.

The ten demonstrations are:

1. Single Replacement Reactions
   • Aluminum foil is placed in a solution of copper(II) chloride. The solution heats up as the aluminum foil “dissolves” and a dark, reddish-brown precipitate forms.
   • A copper wire coil is suspended in a silver nitrate solution. The coil develops silver “needles” and the solution becomes blue as silver ions are replaced by copper ions.
2. Double Replacement Reactions
   • Solutions of blue-green copper(II) chloride and colorless sodium phosphate are mixed. The color of the resulting solution fades to pale blue as a bright turquoise solid drops out of the reaction mixture.
   • Colorless solutions of sodium chloride and silver nitrate are combined to produce an “avalanche” of snow-white precipitate.
   • Milk of magnesia (MOM) is mixed with universal indicator and a rainbow of colors is observed as the magnesium hydroxide in the milk of magnesia is neutralized upon addition of simulated stomach acid.
3. Combination Reactions
   • Water is placed in an aluminum pan and solid calcium oxide is added. The two chemicals combine to form calcium hydroxide and the heat liberated is great enough to fry an egg!
   • Burning steel wool provides a dramatic demonstration of the rapid oxidation of metals. This exothermic combination reaction of iron metal and oxygen can also be classified as a combustion reaction.
4. Decomposition Reactions
   • A Petri dish filled with a salt solution containing universal indicator is placed on an overhead projector. An electric current is passed through the solution. Water is decomposed at both the cathode and anode, leading to pH changes that create a rainbow of colors in the solution.
   • A sample of copper(II) carbonate is heated, resulting in a color change in the copper(II) carbonate. An indicator solution in a wine airlock, placed over the flask, changes colors as the gas produced in the decomposition reaction bubbles through it.
5. Combustion Reaction
   • Add a little isopropyl alcohol to a 2-liter soda bottle, ignite the vapors and “whoosh!” The combination of isopropyl alcohol and oxygen produces a rush of gases and a spectacular blue flame.

The demonstrations in this kit may be presented in a variety of ways. The demonstrations can be done randomly for student assessment or the demonstrations can be done in order of reaction type as a review. Student data sheets, along with a table summarizing the five classes of reactions, are included as an optional assessment tool for the instructor.

 

Concepts

• Single replacement reaction
• Oxidation–reduction
• Double replacement reaction
• Precipitation reaction
• Acid–base neutralization
• Combination reaction
• Exothermic reaction
• Combustion reaction
• Decomposition reaction
• Gas producing reaction

Materials Included In Kit

Additional Materials Required

Background

Demonstration 1: Single Replacement Reaction—Aluminum and Copper(II) Chloride
Watch aluminum foil disappear as it is added to a green-blue solution of copper(II) chloride. Observe color changes, production of a gas, formation of solid copper, and a drastic change in temperature.

Demonstration 2. Single Replacement Reaction—Copper and Silver Nitrate
The reaction of copper wire with aqueous silver nitrate shows chemistry in action—delicate silver crystals grow on the wire surface and the color of copper(II) ions gradually appears in solution in this example of a single replacement, oxidation–reduction reaction.

Demonstration 3. Double Replacement Reaction—Copper(II) Chloride and Sodium Phosphate
A blue-green solution of copper(II) chloride is combined with a colorless solution of sodium phosphate. The initial blue color of the combined solution fades and the final products consist of a turquoise solid and a pale blue liquid in this example of a double replacement, precipitation reaction.

Demonstration 4. Double Replacement Reaction—Sodium Chloride and Silver Nitrate
Sodium chloride reacts with silver nitrate via a double replacement reaction to form aqueous sodium nitrate and insoluble silver chloride. The driving force for the reaction is the formation of insoluble silver chloride, which precipitates as a white solid.

Demonstration 5. Double Replacement Reaction—Magnesium Hydroxide and Hydrochloric Acid
Mix milk of magnesia (MOM) with universal indicator and observe the dramatic rainbow of colors as the antacid dissolves in the simulated stomach acid! This is a great demonstration to teach concepts of acid–base neutralization.

Demonstration 6. Combination Reaction—Calcium Oxide and Water
When water is added to calcium oxide, the resulting combination reaction produces calcium hydroxide and enough heat to fry an egg!

Demonstration 7. Combination Reaction—Iron and Oxygen
Burning steel wool provides a dramatic example of the rapid oxidation of metals. The exothermic combination reaction of iron metal with oxygen gas can also be classified as a combustion reaction.

Demonstration 8. Decomposition Reaction—Electrolytic Decomposition of Water
Demonstrate the decomposition reaction of water to hydrogen gas and oxygen gas by electrolysis in a very colorful and dramatic way on an overhead projector.

Demonstration 9. Decomposition Reaction—Copper(II) Carbonate
Challenge students to deduce the reactions responsible for the dramatic color changes observed when a sample of copper carbonate is heated.

Demonstration 10. Combustion Reaction—Isopropyl Alcohol and Oxygen
Wow your students with a whoosh! Students will love to see the blue alcohol flame shoot out the mouth of the bottle in this highly exothermic combustion reaction.

Materials

Safety Precautions

Copper(II) chloride solution is toxic by ingestion. Small volumes of hydrogen gas are produced from the reaction. Hydrogen is a highly flammable gas; avoid contact with flames, sparks and other sources of ignition. Silver nitrate solution is mildly toxic by ingestion and it will stain skin and clothes. Avoid contact of all chemicals with eyes, skin and clothing. Sodium phosphate solution is each moderately toxic by ingestion. Avoid contact of all chemicals with eyes and skin. Milk of magnesia is intended for laboratory use only; it has been stored with other non–food-grade laboratory chemicals and is not meant for human consumption. Hydrochloric acid solution is toxic by ingestion and inhalation and is corrosive to skin and eyes. Universal indicator solution is an alcohol-based flammable solution. Calcium oxide is a corrosive solid and a severe body tissue irritant. Avoid all body tissue contact. Reaction of calcium oxide and water will produce large amounts of heat—skin burns are possible. A lump of calcium oxide may disintegrate violently and splatter when water is added. Demonstration 6 should be a teacher demonstration only. Do not allow students to perform this procedure. Food-grade items that are brought into the lab are considered laboratory chemicals and should not be consumed. Do not eat the egg after it is cooked. Acetone is flammable, a dangerous fire risk and toxic by inhalation and ingestion. Use only with proper ventilation and keep away from any open flame or ignition source. Remove all flammable material from the demonstration area. Copper(II) carbonate is slightly toxic by ingestion and inhalation. 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. 

Please read all safety precautions before proceeding with the Combustion Reaction—Isopropyl Alcohol and Oxygen demonstration.

• Do not perform demo under sprinkler head or smoke detector.
• Isopropyl alcohol is a flammable liquid and a fire hazard. It is slightly toxic by ingestion and inhalation. Use in a well-ventilated room.
• Always recap the alcohol bottle and move it far from the demonstration area. Never leave an open bottle of alcohol in the vicinity of the demonstration.
• A safety shield is highly recommended for explosions. Even the mildest explosion creates some chance of shattering and flying objects. Protective eyewear must be worn by the demonstrator as well as by anyone viewing the demo.
• Never perform alcohol explosions in glass bottles. The large quantities of gases (H₂O and CO₂) produced during the rapid combustion will easily shatter a glass container. Use a bottle made of PETE (polyethylene teraphthalate).
• The alcohol in the bottle may flash back or continue to burn in the bottle. Be prepared to cover the bottle with a fire blanket if this happens.
• Always pour out any excess liquid isopropyl alcohol from the bottle before igniting the vapor. If any liquid isopropyl alcohol is left, it will increase the amount of gaseous afterburning. The liquid could also ignite, which may cause the plastic bottle to melt.
• Always keep a lid or some sort of cover handy, which can be placed over the mouth of the bottle to extinguish the flame if it continues so long as to begin melting the bottle. Excess alcohol on the outside of the bottle should be wiped off in order to avoid its igniting and softening the plastic bottle.
• Never, ever use a pure oxygen environment as the potential for an extremely violent and deadly explosion is possible.
• Never use methyl alcohol for this demonstration. The high volatility of methyl alcohol means that it has the potential for the most violent combustion of any alcohol.
• Replace the plastic bottle after every demonstration. The heat of reaction will cause the bottle to contract and deform.
• Always wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron when performing this demonstration.
• Please consult current Safety Data Sheets for additional safety information on isopropyl alcohol.

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. Allow the solid material in the beakers from Demonstration 1 to settle. Decant the copper(II) chloride solution down the drain according to Flinn Suggested Disposal Method #26b. Dispose of the solid copper and leftover aluminum foil in the solid waste according to Flinn Suggested Disposal Method #26a. The copper wire and the plated silver from Demonstration 2 may be carefully removed from the solution and disposed of according to Flinn Suggested Disposal Method #26a. The solution may be disposed of according to Flinn Suggested Disposal Method #26b. The product mixture from Demonstration 3 can be filtered, with the copper(II) phosphate solid disposed of according to Flinn Suggested Disposal Method #26a and the remaining sodium chloride filtrate may be disposed of according to Flinn Suggested Disposal Method #26b. The product mixture from Demonstration 4 can be filtered, with the silver chloride solid disposed of according to Flinn Suggested Disposal Method #26a and the sodium chloride filtrate disposed of according to Flinn Suggested Disposal Method #26b. Neutralize the final solution from Demonstration 5 according to Flinn Suggested Disposal Method #24b. Excess milk of magnesia can be disposed of according to Flinn Suggested Disposal Method #26a. The solid calcium hydroxide produced in Demonstration 6 may be neutralized with hydrochloric acid and flushed down the drain with excess water according to Flinn Suggested Disposal Method #10. The steel wool from Demonstration 7 can be discarded in the trash according to Flinn Suggested Disposal Method #26a. Excess acetone can be disposed of according to Flinn Suggested Disposal Method #18a. All materials from Demonstration 8 may be disposed of according to Flinn Suggested Disposal Method #26b. The powdered contents of the flask and any unused carbonate from Demonstration 9 may be disposed of according to Flinn Suggested Disposal Method #26a. The airlock solution may be disposed of according to Flinn Suggested Disposal Method #26b. Excess isopropyl alcohol from Demonstration 10 may be disposed of by allowing it to evaporate in a fume hood according to Flinn Suggested Disposal Method #18a. The bottle may be disposed of according to Flinn Suggested Disposal Methos #26a.

Prelab Preparation

Demonstration 7. Combination Reaction—Iron and Oxygen
Place steel wool in a 150-mL beaker, cover it with acetone, and allow the steel wool to soak in the acetone for 15–20 minutes. Remove the steel wool and allow it to dry inside a fume hood. The acetone removes the oil or plastic coating that is usually present on steel wool to prevent rust. Caution: Be sure all the acetone has evaporated from the steel wool before igniting the steel wool. Pour the acetone back into its bottle, cap the bottle, and remove it from the demonstration area before igniting the steel wool.
 
Demonstration 9. Decomposition Reaction—Copper(II) Carbonate
Insert the stem of the wine airlock into the size #6 1-hole rubber stopper until the stem is completely through the stopper. Remove the cover from the top of the airlock.

Demonstration 10. Combustion Reaction—Isopropyl Alcohol and Oxygen
Before each demonstration, inspect the plastic 2-L bottle for grazing, frosting, cracking, or any small flaws. Replace the bottle if it shows signs of fatigue.

Procedure

Demonstration 1: Single Replacement Reaction—Aluminum and Copper(II) Chloride

1. Place a 600-mL borosilicate beaker (or a 500-mL borosilicate graduated cylinder) on the demonstration table.
2. Use a 250-mL graduated cylinder to measure 140 mL of 1.0 M CuCl₂ solution. Pour this solution into the 600-mL beaker.
3. Measure and add 140 mL of distilled or deionized water to the beaker. The solution is now 0.5 M CuCl₂.
4. Cut a piece of aluminum foil approximately 6" x 12". Loosely roll the foil enough to fit into the beaker. (Note: Do not wad up the foil tightly in a ball as this decreases the surface area and slows any reaction that may occur.)
5. Measure the temperature of the solution before adding the aluminum foil.
6. Place the aluminum foil into the beaker, using a stirring rod to push it down completely into the solution. Measure the temperature of the reaction mixture again. Notice the great increase in temperature, indicating a very exothermic reaction.
7. Have the students record the evidence for a chemical reaction and the properties of the product(s) in the data table.

Demonstration 2. Single Replacement Reaction—Copper and Silver Nitrate

1. Cut a 30-cm piece of copper wire and loosely coil it into the shape shown in Figure 1.
   {13940_Procedure_Figure_1}
2. Add distilled water to the 120-mL mark of a clean, 250-mL beaker.
3. Fill the beaker to the 160-mL mark (add 40 mL) with 0.10 M AgNO₃. Stir the solution with a clean stirring rod.
4. Use a wood splint to suspend the copper wire in the silver nitrate solution, as shown in Figure 1. The copper wire should not be touching the bottom or sides of the beaker.
5. The reaction will take 15–25 minutes before the needles of silver are clearly visible. Set the beaker aside.
6. Instruct the students to check the beaker periodically and record any evidence of a chemical reaction and the properties of the product(s) in the data table.

Demonstration 3. Double Replacement Reaction—Copper(II) Chloride and Sodium Phosphate

1. Use a 100-mL graduated cylinder to transfer 70 mL of the 0.05 M copper(II) chloride solution to the 250-mL graduated cylinder or the hydrometer cylinder.
2. Measure 50 mL of 0.05 M sodium phosphate solution using a clean, 100-mL graduated cylinder.
3. Have the students record the initial colors of the two solutions in the data table.
4. Add 50 mL of sodium phosphate solution to the 250-mL graduated cylinder. Stir the solution to mix.
5. Have the students record the evidence for chemical reaction and the properties of the product(s) in the data table.

Demonstration 4. Double Replacement Reaction—Sodium Chloride and Silver Nitrate

1. Place the 250-mL beaker with a stirring rod on the demonstration table.
2. Use the graduated cylinder to add 100 mL of the 0.1 M sodium chloride solution to the beaker.
3. Add distilled or deionized water to the 175-mL mark of the beaker.
4. Add 25 mL of the 0.1 M silver nitrate solution to the beaker. Stir.
5. Have the students record the evidence of a chemical reaction and the properties of the product(s) in the data table.

Demonstration 5. Double Replacement Reaction—Magnesium Hydroxide and Hydrochloric Acid

1. Measure 20 mL of milk of magnesia using a 25-mL graduated cylinder and pour it into a 1-L beaker.
2. Place the 1-L beaker on a magnetic stir plate. Add a magnetic stir bar to the beaker.
3. Add water and crushed ice (or ice cubes) to give a total volume of approximately 800 mL. Turn on the stir plate to create a vortex in the mixture.
4. Using the Beral-type pipet, add about 4–5 mL (about 2 pipets full) of universal indicator solution. Watch as the white suspension of milk of magnesia turns to a deep purple color. The color indicates that the solution is basic.
5. With a second Beral-type pipet, add 2–3 mL (1 pipet-full) of 3.0 M HCl. The mixture quickly turns red and then passes through the entire range of universal indicator color changes back to purple.
6. Repeat this process, adding HCl one pipet-full at a time, waiting after each addition until the mixture turns back to purple.
7. The process can be repeated a number of times before all of the Mg(OH)₂ dissolves and reacts with the HCl. As more acid is added, the color changes will continue to occur, but the indicator color will not return to purple. When the Mg(OH)₂ has completely reacted, the final solution will be clear and red.
8. Have students record the evidence for a chemical reaction and the properties of the product(s) in the data table.

Demonstration 6. Combination Reaction—Calcium Oxide and Water

1. Measure 100 g of calcium oxide into a large weighing dish.
2. Place one of the aluminum pie pans onto a heat-resistant surface or pad. Using a graduated cylinder, measure and add 40 mL of distilled water to the pan.
3. Spray the second aluminum pan with cooking spray or add a small amount of cooking oil to the pan.
4. Quickly but carefully add the calcium oxide to the water in the first aluminum pan. Distribute the calcium oxide evenly across the pan to get efficient contact between the solid calcium oxide and water. Caution: Splattering may occur.
5. Immediately place the second aluminum pan directly on top of the calcium oxide.
6. Wait about one minute—steam should be evident emanating from the bottom pan, and the cooking oil or spray may begin to sizzle.
7. Quickly break a small egg into the top pan. Caution: The pan will be hot. Wear a heat-resistant glove or use an oven mitt to hold the top pan with one hand while gently using the spatula to check the egg with the other hand.
8. Cook the egg to order! Continue cooking for several minutes—the egg white will continue to cook after the steam has subsided. (The egg white will “fry” but the yolk will remain runny.)
9. Have the students record the evidence for chemical reaction and the properties of the product(s) in the data table.

Demonstration 7. Combination Reaction—Iron and Oxygen

1. Accurately determine the mass of about one gram of steel wool using a balance with 0.01-g precision. Use a plastic weighing dish to mass the steel wool and to capture the reaction products.
2. Unwrap the steel wool, pulling the strands apart as far as possible to get maximum surface exposure.
3. Hang the pulled-apart steel wool from a clamp on a ring stand or hold it with tongs.
4. Ignite the steel wool using a match, lighter or Bunsen burner. The steel wool will burn readily. Some steel wool strands will break off from the bulk. However, if the combustion is done over the plastic weighing dish, most of the reaction product(s) can be captured.
5. After about two minutes, the product should be cool enough to handle. Ask students to predict if the mass of the product will be greater or less than that of the starting steel wool. Return the product to the weighing dish and determine its mass.
6. Have the students record the evidence for a chemical reaction and the properties of the product(s) in the data table.

Demonstration 8. Decomposition Reaction—Electrolytic Decomposition of Water

1. Place the two halves of a Petri dish on the projection stage of an overhead projector.
2. Pour enough sodium chloride/universal indicator solution into each half of the Petri dish to just cover the bottom of each half dish. Adjust the overhead so that the dishes are in clear focus. Each half dish should appear to be a rich, transparent green color.
3. Break a pencil lead in half. Attach the leads to opposite sides of the Petri dish with the alligator clips. Make sure the tip of each lead is submerged in the green solution and the alligator clips remain out of the solution.
4. To start the demonstration, clip the 9-volt battery into the snaps on the battery clip (see Figure 2).
   {13940_Procedure_Figure_2}
5. Let the demonstration proceed for 5–10 minutes and note the changing colors over time. (A purple color will appear at the cathode very quickly. An orange color at the anode will appear more slowly. Over time, the entire spectrum of universal indicator colors will appear.)
6. Have the students record the evidence for chemical reaction and the properties of the product(s) in the data table as the demonstration continues. Discuss the various colors as well as why the “extra” dish was included in the demonstration.

Demonstration 9. Decomposition Reaction—Copper(II) Carbonate

1. Add 5 g of copper(II) carbonate solid to a 250-mL Erlenmeyer flask.
2. Add 10 mL of bromthymol blue indicator solution to the wine airlock.
3. Insert the stopper with the wine airlock into the flask (see Figure 3).
   {13940_Procedure_Figure_3}
4. Set the flask on the hot plate and begin heating the apparatus. As the flask is heated gas, bubbles will be observed flowing through the solution in the airlock.
5. Upon further heating, the color of the bromthymol blue solution in the airlock will change from blue to yellow.
6. The color of the powder in the flask will change from green to black as the decomposition reaction proceeds.
7. Have the students record the evidence for a chemical reaction and the properties of the product(s) in the data table.

Demonstration 10. Combustion Reaction—Isopropyl Alcohol and Oxygen

1. Add about 5 mL of isopropyl alcohol to the 2-L plastic bottle. Do not add more than 8 mL of alcohol. Recap the bottle of alcohol tightly and move it far from the demonstration area.
2. Lay the bottle sideways on a flat surface, allowing the alcohol to flow from base to mouth. Slowly swirl the jug for about 30 seconds, trying to spread alcohol liquid completely over the entire interior surface. This allows the liquid alcohol to volatilize and makes the vapor concentration uniform throughout the bottle. If a lot of liquid alcohol is still visible, swirl the bottle for another 30 seconds.
3. Pour out any excess liquid alcohol into a beaker. Remove the beaker from the area before igniting the vapor in the bottle.
4. Stand the bottle on the lab bench, placing it in the front of the room and behind a safety shield.
5. Dim the lights in the room.
6. Light a match or wood splint that is taped to a meter stick or other long stick.
7. Stand back and, at arm’s length, bring the burning match or wood splint over or slightly down into the mouth of the bottle. (Note: Be sure you are on the safe side of the safety shield as well.)
8. Observe the explosive “whoosh” that results.
9. After the reaction has subsided and all the flames are out, wait for a few minutes until the bottle has cooled slightly. Pour out the water droplets from the bottle into a 10-mL graduated cylinder using a small funnel. As much as 2 mL of water may result, showing that water is one of the products of the combustion of alcohol.
10. Have the students record the evidence for chemical reaction and the properties of the product(s) in the data table.
11. Repeating the Demonstration: The demonstration cannot be repeated with the same bottle. Use a new 2-L plastic bottle to repeat the demonstration. To repeat the demonstration for the same class or another class, follow steps 1–9.

Student Worksheet PDF

13940_Student1.pdf

Teacher Tips

• This kit contains enough chemicals to perform the Single Replacement Reaction—Aluminum and Copper(II) Chloride demonstration seven times using the amounts indicated in the procedure. You may opt to scale the reaction up or down proportionally, depending on class size.
• The reaction temperature increases from room temperature (25 °C) to nearly 60 °C when aluminum is added to copper(II) chloride. Be sure to perform the demonstration in borosilicate glassware.
• The use of a digital thermometer allows for quick temperature measurements.
• A more dramatic demonstration display can be achieved by performing the demonstration in a large 500-mL borosilicate graduated cylinder or hydrometer cylinder. In this case, loosely coil up the foil into a tube-like piece and drop it into the cylinder containing the CuCl₂ solution. The foil will begin to slowly rise up in the cylinder from the gas bubbles attaching to the foil piece. Most of the reaction will occur at the top of the solution (which will turn gray or colorless) and there will be unreacted green-blue copper solution at the bottom of the cylinder. A long stirring rod will be useful in pushing the foil down to the bottom.
• Add more aluminum foil to get the final solution to colorless.
• Exact amounts of reactants are not critical.
• This kit contains enough chemicals to perform the Single Replacement Reaction—Copper and Silver Nitrate demonstration seven times using the amounts indicated in the procedure.
• The copper wire may be surface cleaned with a paper towel and reused for each demonstration.
• The crystals of silver that form will range from white to gray in appearance.
• Three to four drops of 3 M nitric acid may be added to the reaction mixture to initiate the reaction and speed up the demonstration.
• This kit contains enough chemicals to perform the Double Replacement Reaction—Copper(II) Chloride and Sodium Phosphate demonstration seven times using the amounts indicated in the procedure.
• The final solution color depends on the concentration of unreacted copper(II) ions. The fewer the ions the paler the final solution color.
• Add more 0.05 M sodium phosphate to completely clear the solution of copper(II) ions.
• If using the same graduated cylinder to add both reagents, be sure to thoroughly rinse the cylinder with distilled or deionized water.
• This kit contains enough chemicals to perform the Double Replacement Reaction—Sodium Chloride and Silver Nitrate demonstration seven times using the amounts indicated in the procedure.
• The precipitated silver chloride will remain suspended initially, then gradually settle out over time.
• To reclaim the silver, use Flinn Suggested Disposal Method #11.
• If using the same graduated cylinder, thoroughly rinse the cylinder with distilled or deionized water.
• This kit contains enough materials to perform the Double Replacement Reaction—Magnesium Hydroxide and Hydrochloric Acid demonstration as written seven times: 150 mL of milk of magnesia, 250 mL of 3.0 M hydrochloric acid solution, 50 mL of universal indicator solution, and 15 Beral-type pipets.
• If a 1-L beaker is not available, use a 600-mL or 400-mL beaker. Adjust chemical amounts accordingly. Note: The actual milk of magnesia concentration does not have to be exact in order for the demonstration to work.
• If a magnetic stir plate is not available, the mixture can be stirred with a stirring rod.
• The acid used in the demonstration is 3.0 M hydrochloric acid (HCl). Actual stomach acid ranges in concentration from 0.1 to 1.0 M HCl. However, 3.0 M HCl is used in this demonstration in order to limit the total acid volume and allow the reaction to go to completion with a reasonable volume of acid. If desired, dilute the 3.0 M acid to 1.0 M and perform the experiment as written. The volume of acid needed will be three times greater.
• The reaction is performed with ice in order to slow down the color changes so that all the colors in the universal indicator color range can be observed. The reaction may also be performed without the use of ice.
• Consider performing this demonstration at different temperatures—5 °C, 25 °C and 60 °C—to illustrate the effect of temperature on the rate of reaction.
• This kit contains 700 g of CaO, enough to perform the Combination Reaction—Calcium Oxide and Water demonstration at least seven times. The shelf life for calcium oxide is poor—always use fresh calcium oxide for best results.
• Use an oven mitt or hot pad to hold the pans when cooking the egg. Place the aluminum pans on a heat-resistant surface— the bottom pan will get very hot. This reaction generates a lot of heat; use proper care when handling the pans. Steam burns are possible—the maximum temperature of the solid mixture may reach 120 °C.
• This kit contains 10 g of steel wool and 50 mL of acetone, enough to perform the Combination Reaction—Iron and Oxygen demonstration seven times.
• Students will usually say the combusted steel wool will lose mass, like the ash from burning paper. Others may predict no change (conservation of mass). The mass increases due to the addition of oxygen.
• This kit contains enough materials to perform the Decomposition Reaction—Electrolytic Decomposition of Water demonstration at least seven times. The 9-V battery and battery leads with alligator clips are reusable.
• Concepts of pH and electrolysis should be discussed prior to this demonstration. Universal indicator colors, as they relate to pH values, should also be discussed. The Flinn Overhead Color Chart (Catalog No. AP5367) would be a nice visual supplement during this discussion.
• If the sodium chloride/universal indicator solution is not initially green, add a small amount of acid or base to return it to the green color. (See Table 1 in the Discussion section for Demonstration 5.)
• This kit contains enough materials to perform the Decomposition Reaction—Copper(II) Carbonate demonstration seven times using the amounts indicated in the procedure.
• Make sure to use a flask made of borsilicate glass.
• Bromthymol blue is an acid–base indicator that is yellow when the pH is ≤ 6.0, blue when the pH is >7.6, and various shades of green in the transition pH range of 6.0–7.6.
• In addition to recognizing the decomposition of carbonate, the students should be able to answer the following questions.
  • What are the bubbles being produced?
  • Why is the indicator solution changing colors?
  • Why has the powder changed colors?
• Break the seal between the flask and the rubber stopper before removing the Erlenmeyer flask from the hot plate.
• Make sure the bromthymol blue indicator solution starts out blue—if not, add a drop of 0.01 M NaOH.
• Enough isopropyl alcohol (50 mL) is provided to perform the Combustion Reaction—Isopropyl Alcohol and Oxygen demonstration at least 10 times.
• Do not try this demonstration with methyl alcohol. The high volatility of methyl alcohol means that one must be particularly cautious when using methyl alcohol as it has the potential for the most violent combustion and possible rupture of the bottle.
• Depending on how much alcohol vapor is in the bottle, you may have to place the flame slightly inside the lip of the bottle before it ignites.
• The demonstration works best if the alcohol vapor is prepared immediately before the demonstration. If the bottle with the vapor sits for a while, the vapor tends to settle and is harder to light.
• Use a graduated cylinder to measure the volume of water produced by the reaction.
• Flinn sells a larger version of this demonstration, Whoosh Bottle, Catalog No. AP5943.

Sample Data

See Teacher PDF.

Discussion

Demonstration 1: Single Replacement Reaction—Aluminum and Copper(II) Chloride
Aluminum foil reacts with an aqueous solution of copper(II) chloride according to Equation 1. The reaction may be classified as a single replacement, oxidation–reduction reaction.

The oxidation of aluminum metal to aluminum(III) (Al⁰ to Al³⁺) is inferred from the dissolving of the aluminum foil and is represented by the oxidation half-reaction below. The simultaneous reduction of copper(II) or copper(II) ions to copper metal [Cu²⁺(aq) to Cu(s)] will occur and solid copper metal precipitates from solution according to the reduction half reaction that follows. As the copper(II) ions are reduced to copper, the green-blue solution color will fade until the solution is completely colorless—the indication that the reaction is complete and all of the copper(II) ions have been reduced.

It is observed that hydrogen gas is simultaneously released from the reaction when aluminum metal foil is added to copper(II) chloride solution. If the pH of the copper(II) chloride solution is measured, it is found to be slightly acidic. Hence there are free hydrogen ions in solution, which cause the side reaction of hydrogen ions with the aluminum surface to form hydrogen gas and aluminum ions (see Equation 2). Due to the limited concentration of hydrogen ions, this reaction consumes only a small amount of the aluminum.

An interesting note is that the reaction will not work with a copper(II) sulfate, CuSO₄, solution. Chloride ions are needed in solution to penetrate the aluminum oxide coating and expose the aluminum metal to oxidation.

Demonstration 2. Single Replacement Reaction—Copper and Silver Nitrate
The silver ions in solution are reduced by the copper metal in an oxidation–reduction reaction.

The copper(II) ions formed during the reaction will impart a blue color to the solution.

Demonstration 3. Double Replacement Reaction—Copper(II) Chloride and Sodium Phosphate
A solution of sodium phosphate reacts with a copper(II) chloride solution according to Equation 3. The reaction may be classified as a double replacement, precipitation reaction.

Insoluble copper(II) phosphate, Cu₃(PO₄)₂, precipitates out of solution as a turquoise colored solid. The removal of most of the blue-green copper(II) ions, Cu²⁺(aq), from solution reduces the solution color to a very pale green.

Demonstration 4. Double Replacement Reaction—Sodium Chloride and Silver Nitrate
Sodium chloride reacts with silver nitrate to form insoluble silver chloride and sodium nitrate, as shown in Equation 4. The driving force for the reaction is the formation of insoluble silver chloride, which precipitates out of the reaction mixture to give a milky white solution.

Demonstration 5. Double Replacement Reaction—Magnesium Hydroxide and Hydrochloric Acid
The active ingredient in milk of magnesia is magnesium hydroxide, Mg(OH)₂. Magnesium hydroxide forms a suspension in water since it has a very low solubility—0.0009 g/100 mL in cold water and 0.004 g/100 mL in hot water.

The initial reaction mixture is basic due to the small amount of Mg(OH)₂ that goes into solution. The universal indicator gives the mixture a violet color, indicating a pH of about 10 (see Table 1). When hydrochloric acid (the simulated “stomach acid”) is added, the mixture quickly turns red (acidic) because the acid disperses throughout the beaker and then neutralizes the small amount of dissolved Mg(OH)₂.

As more of the Mg(OH)₂ goes into solution, the acid is neutralized and the mixture becomes basic again from the excess Mg(OH)₂(aq) that is present. The addition of universal indicator allows this dissolving process to be observed. During the process, the color of the mixture cycles through the entire universal indicator color range—from red to orange to yellow to green to blue and finally back to violet. By adding more “stomach acid,” the process can be repeated several times before all of the Mg(OH)₂ dissolves and is neutralized. After all of the Mg(OH)₂ has reacted, the final solution will be red (acidic) and clear (no solids).

agnesium hydroxide is classified as a weak base due to its very limited solubility in water. This limited solubility makes it an ideal compound to use in commercial antacids since it slowly dissolves as it neutralizes stomach acid rather than dissolving all at once. The neutralization reaction is the reaction between Mg(OH)2 (a weak base) and HCl (a strong acid). The overall equation for the reaction is shown in Equation 5:

Demonstration 6. Combination Reaction—Calcium Oxide and Water
Calcium oxide is also known as lime or quicklime. Calcium oxide is produced by heating limestone (calcium carbonate) in air. However, calcium oxide readily absorbs and reacts with carbon dioxide and water to form calcium carbonate (CaCO₃) and calcium hydroxide [Ca(OH)₂], respectively.

When water is added to calcium oxide, an exothermic, combination reaction occurs, producing calcium hydroxide and a large amount of heat. Calcium hydroxide is used to treat acidic soils, soften water, and prepare many building materials such as plaster, mortar, and bricks. The solubility of calcium hydroxide in water is very low, about 1.6 g/L. The product of the reaction of CaO and H₂O is thus Ca(OH)₂(s), not Ca(OH)₂(aq).

CaO(s) + H₂O(l) → Ca(OH)₂(s) + heat
ΔH = ΔH_f (products) – ΔH_f(reactants)
ΔH = ΔH_f [Ca(OH)₂(s)] – {ΔH_f [CaO(s)] + ΔH_f [H₂O(l)]}
ΔH = –986.1 kJ/mole – [–635.1 kJ/mole + (–285.8 kJ/mole)] = –65.2 kJ/mole

Demonstration 7. Combination Reaction—Iron and Oxygen
The oxidation of steel wool may be classified as either a combination reaction or a combustion reaction. Many students don’t associate combustion with oxidation. Many elements, particularly metals, will burn. Finely divided metals, especially iron, magnesium, and zinc dusts, are extremely flammable. The key to burning metals is the surface area of the metal and the amount of oxygen available. Airborne metal dusts or fine strands of metals have large surface areas where rapid oxidation can occur. For iron, two combustion reactions are possible (Equations 6 and 7). Ferric oxide (Fe₂O₃) has a rust color, similar to the typical rust color of ferric hydroxide Fe(OH)₃, which is produced when iron is oxidized in the presence of water. Ferrosoferric oxide (Fe₃O₄) is a bluish-gray color and is the main oxide produced when iron burns in air. Note: Fe₃O₄ is a mixed Fe(II)/Fe(III) oxide.

If complete combustion occurs, the net gain of mass should be about 14% (M.W. Fe₃O₄/3Fe = 231.5/167.5). However, due to incomplete combustion and losses due to the uncollected residue, mass increases of around 10% are more common.

Demonstration 8. Decomposition Reaction—Electrolytic Decomposition of Water
When an electric current is passed through an aqueous solution containing an electrolyte (NaCl), the water molecules break apart or decompose into their constituent elements, hydrogen and oxygen. The overall reaction occurs as two separate, independent half-reactions. Reduction of the hydrogen atoms in the water molecules to elemental hydrogen (H₂) occurs at the cathode (–), while oxidation of the oxygen atoms in water to elemental oxygen (O₂) occurs at the anode (+). Each half-reaction is accompanied by the production of OH^– or H⁺ ions as shown:

4e^– + 4H₂O(l) → 2H₂(g) + 4OH^–(aq) Cathode
2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e^– Anode

At the cathode, the excess OH^– ions will cause the pH to increase, resulting in a color change of the universal indicator solution from green (neutral, pH 7) to purple (basic, pH ≥ 10).

At the anode, the excess H⁺ ions will cause the pH to decrease, resulting in a color change of the universal indicator solution from green to an orange/red color (acidic, pH ≤ 4). The electrolysis half-reactions can also be followed by observing the production of gas bubbles at the cathode (H₂) and anode (O₂).

Universal indicator is an acid–base indicator that is different colors at different pH values. All colors will be visible in the Petri dish as electrolysis proceeds and as the pH conditions continually change due to diffusion and neutralization.

Demonstration 9. Decomposition Reaction—Copper(II) Carbonate
When a compound undergoes decomposition, it breaks apart into two or more simpler substances. Copper(II) carbonate will decompose to form copper(II) oxide, carbon dioxide and water upon heating. Note: The initial substance in the flask is “basic” copper carbonate, Cu₂CO₃(OH)₂. Copper(II) carbonate does not exist as CuCO₃.

In this demonstration, an acid–base indicator (bromthymol blue) solution, held in a winemaker’s airlock, changes color as the CO₂ produced during the decomposition of copper(II) carbonate passes through the airlock. Furthermore, the green copper(II) carbonate powder changes to black copper(II) oxide.

Initially, as the apparatus is heated, the air in the flask will be forced through the bromthymol blue solution in the airlock due to the air expanding upon heating. This will be observed as bubbles moving through the airlock.

As the apparatus is further heated, the bromthymol blue solution color changes from blue to yellow to indicate the presence of an acid. As the carbon dioxide gas formed from the decomposition reaction dissolves in the bromthymol blue aqueous solution, it forms carbonic acid (Equation 9). This weak acid dissociates to produce an acidic solution (Equation 10).

Demonstration 10. Combustion Reaction—Isopropyl Alcohol and Oxygen
Low-boiling alcohols vaporize readily. When isopropyl alcohol is placed in a 2-liter PETE plastic bottle, it forms a volatile mixture with the air. A simple match held by the mouth of the bottle provides the activation energy needed for the combustion of the alcohol/air mixture.

Only a small amount of alcohol is used and it quickly vaporizes to a heavier-than-air vapor. The alcohol vapor and air are all that remain in the bottle. Alcohol molecules in the vapor phase are farther apart than in the liquid phase and present far more surface area for reaction; therefore the combustion reaction that occurs is very fast. Since the burning is so rapid and occurs in the confined space of a 2-Liter plastic bottle with a small neck, the sound produced is very interesting, like a “whoosh.” The equation for the combustion reaction of isopropyl alcohol is shown below. One mole of isopropyl alcohol reacts with 4.5 moles of oxygen to produce 3 moles of carbon dioxide and 4 moles of water:

(CH₃)₂CHOH(g) + ⁹⁄₂O₂(g) → 3CO₂(g) + 4H₂O(g)           ΔH = –1,235 kJ/mol

References

Demonstration 1: Single Replacement Reaction—Aluminum and Copper(II) Chloride
Special acknowledgement to the late Cliff Schrader, Summit County ESC, Cuyahoga Falls, OH, for providing Flinn with the idea for this activity. Andy Cherkas, Walter Rohr, and Pat Funk also provided insight and advice on this activity.

Demonstration 5. Double Replacement Reaction—Magnesium Hydroxide and Hydrochloric Acid
Special thanks to Annis Hapkiewicz, Okemos High School, Okemos, MI, and to Penney Sconzo, Westminster School, Atlanta, GA, for separately bringing this demonstration to our attention.

Demonstration 6. Combination Reaction—Calcium Oxide and Water
Special thanks to DeWayne Lieneman (retired), Glenbard South High School, Glen Ellyn, IL, for providing us with the instructions for this activity.

Demonstration 7. Combination Reaction—Iron and Oxygen
Bilash, B. B., Gross, G. R., Koob, J. K. A Demo A Day; Flinn Scientific: Batavia, IL 1995; p 201.

Demonstration 8. Decomposition Reaction—Electrolytic Decomposition of Water
Special thanks to Mike Shaw, West Stokes High School, King, NC, for this demonstration idea.

Demonstration 9. Decomposition Reaction—Copper(II) Carbonate
Special thanks to Borislaw Bilash II, chemistry teacher, Pascack Valley High School, Hillsdale, NJ, for providing Flinn Scientific with the instructions for this demonstration.

Demonstration 10. Combustion Reaction—Isopropyl Alcohol and Oxygen
Flinn Scientific would like to thank John Fortman, Dept. of Chemistry, Wright State University, Dayton, OH, for all of his research, safety notes, and variations on this classic demonstration. Lee Marek, (retired), Naperville, IL, and Bill Deese have also popularized this demonstration.