Teacher Notes

Introduction to Chemical Reactions

Student Laboratory Kit

Materials Included In Kit

Copper(II) carbonate, CuCO3, 40 g
Copper(II) nitrate, Cu(NO3)2, 0.1M, 75 mL
Copper wire, Cu
Ethanol, C2H5OH, 70%, 50 mL
Hydrochloric acid, HCl, 3 M, 200 mL
Magnesium ribbon, Mg
Silver nitrate, AgNO3, 0.1 M, 75 mL
Silver nitrate, AgNO3, 0.5 M, 100 mL
Sodium chloride, NaCl, 0.1 M, 75 mL
Sodium phosphate, Na3PO4, 0.1 M, 100 mL
Zinc metal, Zn, 60 g
Pipets, 50
Sand paper, 9" x 11"
Wood splint, 30

Additional Materials Required

Matches or butane safety lighter
Test tubes, 6
Test tube holder
Test tube rack
Tongs
Watch glass, borosilicate

Prelab Preparation

Sand paper

  1. Cut the sand paper into 16 pieces.
Copper Wire
  1. Cut 15 3-cm pieces of copper wire.
  2. Cut 15 8-cm pieces of copper wire.
Magnesium Wire
  1. Cut the magnesium ribbon into 1–2 cm pieces.

Safety Precautions

Ethanol is a flammable liquid and dangerous fire risk. It is irritating to the eyes and skin and toxic by ingestion. Hydrochloric acid is toxic and corrosive to eyes and skin tissue. Magnesium ribbon is a flammable solid and burns with an intense flame. Silver nitrate is mildly toxic by ingestion and it will stain skin and clothes. Copper(II) carbonate is slightly toxic by ingestion and inhalation. Copper(II) nitrate solution is slightly toxic by ingesting and irritating to skin, eyes and mucous membranes. Sodium phosphate monobasic (monohydrate) is moderately toxic by ingestion. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Avoid contact of all chemicals with eyes, skin and clothing. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory. Please review 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. The leftover hydrochloric acid may be disposed of by neutralizing with base and then disposing of down the drain with plenty of excess water according to Flinn Suggested Disposal Method #24b. Excess silver nitrate solution may be disposed of according to Flinn Suggested Disposal Method #11. Copper (II) carbonate and copper oxide may be disposed of according to Flinn Suggested Disposal Method #26a. Solid silver chloride may be collected by filtration and may be disposed of according to Flinn Suggested Disposal Method #26a. Ethyl alcohol, copper(II) nitrate solutions, and unused sodium phosphate and sodium chloride solutions may be disposed of according to Flinn Suggested Disposal Method #26b.

Lab Hints

  • Enough materials are provided in this kit for 30 students working in pairs or for 15 groups of students. This laboratory activity can reasonably be completed in two 50-minute class periods. Allow students time to identify the type of reaction and write the balanced chemical equations.
  • Another option is to have groups perform one to two of the reactions and then have various groups present their findings. Students could take pictures to share of their reaction(s).
  • Before allowing students to perform the hydrogen test in lab, demonstrate the test for students. Remind students to hold the test tube securely when performing the test.
  • Remind students to write observations of what the items looked like before reacting, during the reaction and after the reaction goes to completion.
  • Remind students to check labels and use the correct concentration of silver nitrate for the appropriate reaction.
  • Rather than have students weigh out the two grams of copper carbonate, prepare an example test tube for them to use as a reference and have them add approximately the same amount.

Teacher Tips

  • Instead of having each group perform all nine reactions at one lab station, you can set up nine reaction stations and have groups rotate for each reaction.
  • Another fun extension to this activity is to assess students by having them perform one of the types of reactions on their own. For example, set up an additional double replacement reaction in the lab. Give the students the reactants and have them perform the reaction and identify what type of reaction it is as well as writing the balanced chemical equation. Ideally, have 3–4 different reaction types for students to be assessed.

Correlation to Next Generation Science Standards (NGSS)†

Science & Engineering Practices

Asking questions and defining problems
Planning and carrying out investigations
Analyzing and interpreting data
Obtaining, evaluation, and communicating information

Disciplinary Core Ideas

HS-PS1.A: Structure and Properties of Matter
HS-PS1.B: Chemical Reactions

Crosscutting Concepts

Patterns
Energy and matter
Structure and function

Performance Expectations

HS-PS1-2. Construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.
HS-PS1-7. Use mathematical representations to support the claim that atoms, and therefore mass, are conserved during a chemical reaction.

Sample Data

{11586_Data_Table_1}

Recommended Products

Item No. Description
AP9894 Flinn Blended Learning Labs for Chemisty: Introduction to Chemical Reactions

Student Pages

Introduction to Chemical Reactions

Concepts

  • Physical and chemical changes
  • Evidence of a chemical reaction
  • Writing and balancing chemical reactions
  • Types of chemical reactions

Background

A chemical reaction results in the transformation of reactants into new substances (products) as a result of a chemical change. Physical changes do not alter the reactants. Physical changes, such as melting ice or ripping paper, change the shape or state of the item, but not the chemical properties. A chemical change, on the other hand, is a change in the chemical composition of a substance.

Signs of a chemical change include:

  • Release of gas bubbles that are not due to a physical change.
  • Energy is released or absorbed (e.g., temperature changes, light)
  • Formation of a solid precipitate
  • A color change that does not result from dilution.
Check out some of these examples of chemical change with the video.

Some chemical reactions may exhibit only one of these signs while other chemical reactions may reveal several signs of a chemical change. By looking for these changes, you can determine whether a chemical reaction has occurred.

Any chemical change involves the reorganization of the atoms in one or more substances.

The chemical equation for a reaction provides important pieces of information. Physical states of the reactants and the products are displayed using abbreviations. Solids are represented with (s), liquids with (l), gases with (g) and aqueous solutions with (aq) to indicate that the substance is dissolved in water. In addition to states of matter, a reaction also has coefficients and subscripts. Subscripts are used in the chemical formulas of compounds and some elements. Coefficients are used to balance the chemical equations. Here’s a sample chemical reaction:
{11586_Background_Equation_1}

Types of Chemical Reactions

There are many different chemical reactions that can occur and different ways to classify them. In this lab, we will be classifying reactions into five basic categories–synthesis reactions, decomposition reactions, single replacement reactions, double replacement reactions and combustion reactions. As you move further into chemistry, you will learn about other types of reactions and ways to classify them, such as acid–base and redox reactions. For now, we will be focusing on these five classifications.

Synthesis Reactions

In a synthesis reaction, two or more substances react to form one new substance. The general form for a synthesis reaction is shown in Equation 2.
{11586_Background_Equation_2}

The reactants A and B may be either elements or compounds while the newly formed product AB is always a compound composed of the elements in A and B. The product AB has different chemical and physical properties than the reactants A and B. Synthesis reactions include reactions like the corrosion of metals in air or water (Equation 3), the reaction of nonmetal oxides with water to produce an acid (Equation 4) and the reaction of a metallic oxide with water to produce a base (Equation 5).
{11586_Background_Equation_3}

{11586_Background_Equation_4}

{11586_Background_Equation_5}

Decomposition Reactions

A decomposition reaction is a reaction in which a single compound AB is broken down, or decomposed, into two or more products, A and B. The products A and B may be either elements or smaller compounds. The general form for a decomposition reaction is shown in Equation 6.
{11586_Background_Equation_6}

Frequently, decomposition reactions occur only when heat is added to the reactant compound AB. The requirement of heat in a reaction is often denoted with a Δ symbol above the arrow. Examples of decomposition reactions that require heat to proceed include the reactions shown in Equations 7 and 8.
{11586_Background_Equation_7}

{11586_Background_Equation_8}

Electrolysis is another common method for carrying out decomposition reactions. In this case, energy in the form of electricity must be added before the reaction will occur. Such reactions include the electrolysis of water to produce oxygen and hydrogen gases (Equation 9) and the electrolysis of table salt, NaCl, to produce liquid sodium and chlorine gas (Equation 10).
{11586_Background_Equation_9}

{11586_Background_Equation_10}

Single Replacement Reactions

Single replacement reactions involve the replacement of one element in a compound with another element. The general form for a single replacement reaction is shown in Equation 11.
{11586_Background_Equation_11}

Examples of single replacement reactions include the reactions shown in Equations 12 and 13.
{11586_Background_Equation_12}

{11586_Background_Equation_13}

During these reactions, electrons are transferred. Single replacement reactions will not occur in the reverse direction without a battery because, in the reverse reaction, energy is needed to force the electrons to move backward. Metals can even be ranked by their level of reactivity in what is called an activity series. This is done by observing how metal solids react with various aqueous salt solutions.

Double Replacement Reactions

Double replacement reactions involve the exchange of ions between two compounds. The general form for a double replacement reaction is shown in Equation 14.
{11586_Background_Equation_14}

The ionic compounds in a double replacement reaction can be thought of as a pair of partners. In Equation 14, A and B are one set of reactant partners while C and D are another set of reactant partners. When these two compounds react, they exchange partners so that A and D become a new set of partners while B and C do the same.

A double replacement reaction generally occurs between two ionic compounds in aqueous solution and is driven by formation of a product that is released from solution, such as in the formation of a precipitate or a gas. Precipitation reactions occur when two soluble compounds react and exchange partners so one of the resulting products is insoluble. An example of this type of reaction is shown in Equation 15.
{11586_Background_Equation_15}

Combustion Reactions
A combustion reaction is a reaction in which a hydrocarbon, or related compound like an alcohol, reacts with oxygen, producing energy in the form of heat and light. A hydrocarbon is a compound that contains hydrogen and carbon. In the complete combustion of a hydrocarbon with oxygen gas, the products are always carbon dioxide gas and water. The general form for complete combustion between a hydrocarbon and oxygen is shown in Equation 16.
{11586_Background_Equation_16}

Because energy is usually produced in substantial quantities during combustion reactions, many hydrocarbons are burned as fuels. Common examples of fuels include methane, propane, butane, octane and gasoline (a mixture of hydrocarbons). The combustion of propane in a gas barbeque grill, for example, is shown in Equation 17.
{11586_Background_Equation_17}

Materials

Copper(II) carbonate, CuCO3, 2 g
Copper(II) nitrate, Cu(NO3)2, 0.1 M, 4 mL
Copper wire, Cu, 3 cm long Copper wire, Cu, 8 cm long
Ethanol, C2H5OH, 0.5 mL
Hydrochloric acid, HCl, 3 M, 10 mL
Magnesium ribbon, Mg, 1–2 cm piece
Silver nitrate, AgNO3, 0.1 M, 1 mL
Silver nitrate, AgNO3, 0.5 M, 5 mL
Sodium chloride, NaCl, 0.1 M, 4 mL
Sodium phosphate, Na3PO4, 0.1 M, 4 mL
Zinc metal, Zn, 1 piece
Matches or butane safety lighter
Pipets
Sand paper, 1 piece
Test tubes, 6
Test tube holder
Test tube rack
Tongs Watch glass, borosilicate
Wood splint

Safety Precautions

Ethanol is a flammable liquid and dangerous fire risk. It is irritating to the eyes and skin and toxic by ingestion. Hydrochloric acid is toxic and corrosive to eyes and skin tissue. Magnesium ribbon is a flammable solid and burns with an intense flame. Silver nitrate is mildly toxic by ingestion and it will stain skin and clothes. Copper(II) carbonate is slightly toxic by ingestion and inhalation. Copper(II) nitrate solution is slightly toxic by ingesting and irritating to skin, eyes and mucous membranes. Sodium phosphate monobasic (monohydrate) is moderately toxic by ingestion. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Avoid contact of all chemicals with eyes, skin and clothing. Wash hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Procedure

Reaction 1

Caution: Perform this reaction away from all flammable materials.

  1. Place approximately 2 grams of copper(II) carbonate into a test tube.
  2. Using a test tube holder, gently heat the sample for 1–2 minutes over a Bunsen burner.
  3. Record your observations.
Reaction 2

Caution: Perform this reaction away from an open flame. Remove all paper and other items from your lab area before proceeding with the reaction.
  1. Using a pipet, obtain 0.5 mL of ethanol.
  2. At your lab station, draw a thin line of ethanol on the chemical and heat resistant counter tops (or place the 0.5 mL of ethanol onto a watch glass).
  3. Using a match or butane safety lighter, light the ethanol.
  4. Once the reaction is complete, try to light the remaining liquid.
  5. Record your observations.
  6. Clean up your area.
Reaction 3

Caution: Silver nitrate can stain skin and clothing.
  1. Place a clean test tube in a test tube rack. Place approximately 4.0 mL of 0.1 M sodium chloride solution into the test tube.
  2. Obtain 1 mL of 0.1 M silver nitrate solution with a pipet.
  3. Slowly add the silver nitrate solution to the test tube, one drop at a time.
  4. Record your observations.
Reaction 4

Caution: Silver nitrate can stain skin and clothing.
  1. Place a clean test tube in a test tube rack. Place approximately 5.0 mL of 0.5 M silver nitrate solution into the test tube.
  2. Obtain a piece of copper wire about 8 cm in length. Wrap the wire around a pencil so it forms a coil.
  3. Place the coiled copper wire into the test tube of silver nitrate solution. Allow the wire to react for approximately five minutes. DO NOT disturb the test tube or contents.
  4. Record your observations.
Reaction 5
  1. Place a clean test tube in a test tube rack. Place approximately 4.0 mL of 0.1 M sodium phosphate solution into the test tube.
  2. Obtain 1 mL of 0.1 M copper(II) nitrate solution with a pipet.
  3. Slowly add the copper(II) nitrate solution to the test tube, one drop at a time.
  4. Record your observations.
Reaction 6

Caution: Perform this reaction away from all flammable materials.
  1. Obtain a three centimeter piece of copper wire.
  2. Using sand paper, clean the piece of copper wire until it is shiny.
  3. Using tongs, hold the wire in the Bunsen burner flame.
  4. Record your observations before and after the copper wire is heated.
Reactions 7 and 8
  1. Place 5–10 mL of 3 M HCl, in a clean test tube.
  2. Add a small piece of zinc metal in the test tube.
  3. Using a second, clean, inverted test tube, collect the gas that is released.
  4. Record your observations.
  5. Place a lit, wooden splint inside the inverted test tube.
  6. Record your observations.
Reaction 9

Caution: Perform this reaction away from all flammable materials.
  1. Obtain a 1–2 cm strip of magnesium metal from your instructor.
  2. Light the laboratory burner.
  3. Hold the piece of magnesium metal with a pair of tongs. Place the ribbon in the burner flame and allow it to burn. DO NOT LOOK DIRECTLY AT THE BURNING MAGNESIUM! The bright light emitted by the burning magnesium ribbon can damage your eyes. Observe by looking slightly to one side and using peripheral vision.
  4. When the reaction is complete, place the remains in a clean watch glass. Turn off the burner.
  5. Record your observations.
  6. Consult your instructor for appropriate disposal procedures.

Student Worksheet PDF

11586_Student1.pdf

†Next Generation Science Standards and NGSS are registered trademarks of Achieve. Neither Achieve nor the lead states and partners that developed the Next Generation Science Standards were involved in the production of this product, and do not endorse it.