Materials Included In Kit

Additional Materials Required

Prelab Preparation

1. Transfer the 12 “unknown” solids into unmarked bottles and label them 1–12. Students can obtain their samples from these bottles. We recommend that half the groups test unknowns 1–6 and the other half test unknowns 7–12. Label the container with potassium nitrate “Do not heat in burner flame.”
2. If you wish, you may transfer approximately 3 g of each unknown solid into four separate capped vials, so there are four vials of each unknown (a total of 48 vials). Use a marker to label the different solids from 1 to 12. Separate the unknowns into eight sets of six vials each. Four of the sets should consist of unknowns 1–6, and the other four sets should consist of unknowns 7–12. Each lab group would get one set of vials to test.

Safety Precautions

Hexane and ethyl alcohol are flammable organic solvents and dangerous fire risks. Keep away from flames, heat and other sources of ignition. Cap the solvent bottles and work with hexane and ethyl alcohol in a fume hood or designated work area. Addition of a denaturant makes ethyl alcohol poisonous; it cannot be made nonpoisonous. Copper(II) sulfate is a skin and respiratory tract irritant and is toxic by ingestion. Graphite powder is a fire and inhalation risk. Iron powder is a serious fire risk. Potassium nitrate is a strong oxidant and a fire and explosion risk when heated or in contact with organic material. It is also a skin irritant. Salicylic acid is moderately toxic by ingestion. Dilute sodium hydroxide and acid solutions are irritating to skin and eyes. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Remind students to wash their hands thoroughly with soap and water before leaving the lab. 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. Sodium hydroxide solution may be neutralized with acid according to Flinn Suggested Disposal Method #10. Hydrochloric acid may be neutralized with base according to Flinn Suggested Disposal Method #24b. The remaining solid samples may be stored for future use or placed in the trash according to Flinn Suggested Disposal Method #26a. Hexane solutions or mixtures require licensed hazardous waste disposal according to Flinn Suggested Disposal Method #18b.

Lab Hints

• If students are unfamiliar with the testing procedures for typical properties, give them a copy of the table in the Teacher PDF.
• Common solids with a wide range of physical properties were deliberately chosen for this study. There is enough overlap to be able to identify patterns in the relationship between the properties of a material and its structure. The challenge in this experiment comes as students try to use their observations to “see inside” the world of atoms and bonds.
• Many common solids may be used. Any metal may be used and many different ionic compounds may be substituted for those in the kit. Suitable nonpolar organic solids that may be used include stearic acid or lauric acid.
• Low-voltage conductivity meters are available from Flinn Scientific (Catalog No. AP1493) for individual student use. The copper wire electrodes are about 2 cm long and are easily inserted into the wells on a microscale reaction plate. Two LEDs make it possible to compare the conductivity of strong versus weak electrolytes. The green LED requires more voltage than the red LED. A weak electrolyte will cause only the red LED to glow. A strong electrolyte will cause both the red and green LEDs to glow. Because the meter uses only a 9-volt battery, the conductivity tester is convenient, portable and safe. Conductivity tests may also be done using conductivity sensors with a computer interface system. Using a conventional 110 -V “lightbulb-type” conductivity tester will require larger samples.
• The recommended organic solvent is hexanes—a mixture of n-hexane and other C₆H₁₄ isomers. Remind students not to use flammable organic solvents around or near flames, heat or other sources of ignition.
• Dodecyl alcohol has a very low melting point of 24 °C. Therefore, in warm environments, it is very likely that this chemical will be in a liquid state. If it cools below 24 °C, it will reform as a solid.
• See the experiment “It’s in Their Nature” in Solubility and Solutions, Volume 12 in the Flinn ChemTopic^™ Labs series, for a detailed investigation into the solubility of ionic, polar and nonpolar compounds in a variety of solvents. Students classify compounds and learn about the different types of attractive forces that exist between molecules.
• The labels for the 12 unknown solids are at the end of the Instructor’s Notes. The labels provide additional information about the properties of the solids, such as solubility, chemical formula and hazard information. Hand these out after students have completed the lab activity.

Further Extensions

Opportunities for Undergraduate Research
The chemistry department suspects vandalism has occurred in the stockroom. An order of chemicals arrived and was placed in the storeroom over the weekend. When the lab manager returned on Monday, the labels on six identically sized bottles appeared misaligned, as if they had been switched. You are asked to use your flow chart to verify or refute the identity of the solids in the six bottles.

Answers to Prelab Questions

1. Considering the data in the above table, explain the following observations based on the type of chemical bonding and intermolecular forces between atoms, molecules or ions in the solid state.
   1. Iodine has a much lower melting point than sodium iodide.

      The forces holding nonpolar molecules, such as I₂, together are London dispersion forces and induced dipole–induced dipole attractions, which are very weak. Those holding ionic compounds together, such as NaI, are strong attractive forces between positive and negative ions.

   2. Iodine is slightly soluble in water, while lactose is very soluble.

      Water, being a polar covalent molecule, dissolves the polar covalent lactose to a much greater extent than with nonpolar covalent iodine. Dipole–dipole interactions and hydrogen bonds between the water molecules and —OH groups in lactose help stabilize the solution.

2. Predict the properties of the following solids.
   {14002_PreLabAnswers_Table_2}

Sample Data

Introductory Activity

Test Results for Unknowns Samples

†Graphite and silicon are semiconductors and do not fit well within the general classification scheme. The solids conduct electricity, but are classified as covalent solids. They can be distiguished from metals based on their lack of luster and ductility.

Answers to Questions

Guided-Inquiry Design and Procedure Questions

1. List the general physical properties associated with each type of bonding in a solid.
   1. Does the solid conduct electricity?

      Yes–Solid is metallic.
      No–Solid may be any of the other three.

   2. Does the solid melt below 100 °C?

      Yes–Solid is nonpolar covalent.
      No–Solid may be any of the other three.

   3. Does the solid dissolve in hexane?

      Yes–Solid is nonpolar covalent.
      No–Solid may be any of the other three.

   4. If the solid dissolves in water, does it conduct electricity?

      Yes–Solid is ionic.
      No–Solid is polar covalent. Nonpolar covalent and metallic substances generally do not dissolve in water.

2. Create a flowchart that can be used to characterize an unknown solid as ionic, polar covalent, nonpolar covalent or metallic.
   {14002_Answers_Figure_1}
   *If solid does not dissolve, test its chemical reaction with dilute HCl. Some ionic compounds, such as carbonates, are only slightly soluble in water but they may react with HCl to produce a gas.
3. A white solid can be identified as polar covalent based on the following physical properties:

   Soluble in water but solution does not conduct (or exhibits slightly) electricity; melts in burner flame (100–500 °C); solid does not conduct electricity.

Post-Laboratory Review Questions

1. Covalent bonds may be classified as polar or nonpolar based on the difference in electronegativity between two atoms. Look up electronegativity values in your textbook:
   1. Why are C—H bonds considered nonpolar?

      The electronegativity values of carbon and hydrogen are similar (2.1 and 2.5, respectively). Both atoms in a C—H bond have similar attractions for the bonding electrons and the bond is nonpolar.

   2. Which is more polar, an O—H or N—H bond? Explain.

      The electronegativity difference between O and H (3.5–2.1) is greater than that between N and H (3.0–2.1). An O—H bond is more polar than an N—H bond.

2. To convert the following compounds from a solid to a liquid, what types of intermolecular forces must be overcome?
   1. I₂(s) → I₂(l) Induced dipole–induced dipole
   2. H₂O(s) → H₂O(l) Dipole–dipole; hydrogen bonding
   3. NaI(s) → NaI(l) Ionic bonding
   4. C₁₆H₃₂(s) → C₁₆H₃₂(l) Induced dipole–induced dipole
3. In order for a substance to conduct electricity, it must have free-moving charged particles.
   1. Explain the conductivity results observed for ionic compounds in the solid state and in aqueous solution.

      Sodium chloride does not conduct electricity in the solid state. It has charged particles (ions) but the ions are "locked" into position in the crystal structure and are not able to move freely. A solution of sodium chloride in water does conduct electricity because the ions are no longer fixed into position. (The solute particles in a liquid are able to move freely.)

   2. Would you expect molten sodium chloride to conduct electricity? Why or why not?

      Molten sodium chloride should conduct electricity because the particles in a liquid are able to move freely.

   3. Use the model of metallic bonding described in the Background section to explain why metals conduct electricity.

      Metals conduct electricity because the valence electrons of the metal are not “attached” to any one metal atom. The electrons are delocalized among all of the metal cations in the crystal structure and are able to move freely throughout the crystal.

Teacher Handouts

14002_Teacher1.pdf

References

Solids and Liquids, Volume 5, Flinn ChemTopic^™ Labs, Cesa, I., Editor; Flinn Scientific, Inc., Batavia IL (2004).

Introduction

Looking for patterns in the properties of solids can help us understand how and why atoms join together to form compounds. What kinds of forces hold atoms together? How do these forces influence the properties of materials? Use your knowledge of the relationship between chemical bonding type and the properties of substances to determine the identity of mystery solids.

Concepts

• Chemical bonds
• Ionic bonding
• Physical and chemical properties
• Covalent bonding
• Metallic bonding
• Electronegativity

Background

Groups of atoms are held together by attractive forces that we call chemical bonds. The origin of chemical bonds is reflected in the relationship between force and energy in the physical world. Think about the force of gravity—in order to overcome the force of attraction between an object and the Earth, we have to supply energy. Whether we climb a mountain or throw a ball high into the air, we have to supply energy. Similarly, in order to break a bond between two atoms, energy must be added to the system, usually in the form of heat, light or electricity. The opposite is also true: whenever a bond is formed, energy is released.

The term ionic bonding describes attractive forces between oppositely charged ions in an ionic compound. An ionic compound is formed when a metal reacts with a nonmetal to form positively charged cations and negatively charged anions, respectively. The oppositely charged ions are arranged in a tightly packed, extended three-dimensional structure called a crystal lattice (see Figure 1). The net attractive forces between oppositely charged ions in the crystal structure are called ionic bonds.

Covalent bonding represents another type of attractive force between atoms. Covalent bonds are defined as the net attractive forces resulting from pairs of electrons that are shared between atoms (the shared electrons are attracted to the nuclei of both atoms in the bond). A group of atoms held together by covalent bonds is called a molecule. Atoms may share one, two or three pairs of electrons between them to form single, double and triple bonds, respectively.

Substances held together by covalent bonds are usually divided into two groups based on whether individual (distinct) molecules exist or not. In a molecular solid, individual molecules in the solid state are attracted to each other by relatively weak intermolecular forces between the molecules. Covalent-network solids, on the other hand, consist of atoms forming covalent bonds with each other in all directions. The result is an almost infinite network of strong covalent bonds—there are no individual molecules.

Covalent bonds may be classified as polar or nonpolar. The element chlorine, for example, exists as a diatomic molecule, Cl₂. The two chlorine atoms are held together by a single covalent bond, with the two electrons in the bond shared between the two identical chlorine atoms. This type of bond is called a nonpolar covalent bond. The compound hydrogen chloride (HCl) consists of a hydrogen atom and a chlorine atom that also share a pair of electrons between them. Because the two atoms are different, however, the electrons in the bond are not equally shared between the atoms. Chlorine has a greater electronegativity than hydrogen—it attracts the bonding electrons more strongly than hydrogen. The covalent bond between hydrogen and chlorine is an example of a polar bond. The distribution of bonding electrons in a nonpolar versus polar bond is shown in Figure 2. Notice that the chlorine atom in HCl has a partial negative charge (δ^–) while the hydrogen atom has a partial positive charge (δ⁺). The special properties of metals compared to nonmetals reflect their unique structure and bonding. Metals typically have a small number of valence electrons available for bonding. The valence electrons appear to be free to move among all of the metal atoms, some of which must exist or act as positively charged cations. Metallic bonding describes the attractive forces that exist between closely packed metal cations and free-floating valence electrons in an extended three-dimensional structure.

Experiment Overview

The purpose of this advanced inquiry lab is to identify twelve unknown solids based on systematic testing of their physical and chemical properties. The lab begins with an introductory activity to select measurable properties that will help identify the type of bonding in a solid. Given four solids representing the four types of chemical bonds—ionic, polar covalent, nonpolar covalent and metallic—students review the properties of each solid with a minimum of four tests. The results provide a basis for a guided-inquiry design of a flow chart procedure to distinguish and identify twelve unknown solids.

Materials

Prelab Questions

1. Considering the data in the above table, explain the following observations based on the type of chemical bonding and intermolecular forces between atoms, molecules or ions in the solid state.
   1. Iodine has a much lower melting point than sodium iodide.
   2. Iodine is slightly soluble in water, while lactose is very soluble.
2. Predict the properties of the following solids.
   {14002_PreLab_Table_2}

Safety Precautions

Hexane and ethyl alcohol are flammable organic solvents and dangerous fire risks. Keep away from flames, heat and other sources of ignition. Cap the solvent bottles and work with hexane and ethyl alcohol in a fume hood or designated work area. Addition of a denaturant makes ethyl alcohol poisonous; it cannot be made nonpoisonous. Copper(II) sulfate is a skin and respiratory tract irritant and is toxic by ingestion. Graphite powder is a fire and inhalation risk. Dilute sodium hydroxide and acid solutions are irritating to skin and eyes. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the lab. Please follow all laboratory safety guidelines.

Procedure

Introductory Activity

Identifying Properties of Chemical Bonds

1. Four representative chemicals are provided for preliminary testing to identify physical and chemical properties that can be used for development of a qualitative analysis scheme. The chemicals and the type of bonding in each are:

   Copper(II) sulfate—ionic bonding
   Paraffin wax—nonpolar covalent bonding
   Dextrose—polar covalent bonding
   Zinc—metallic bonding

2. Observe the color and appearance of each solid and perform the following qualitative tests on each: (a) solubility in water, solubility in hexane and solubility in alcohol; (b) high or low melting point; (c) conductivity of solid and conductivity of aqueous solution; (d) pH of solution; (e) reaction with acid (0.1 M HCl) and reaction with base (0.1 M NaOH).
3. Consult the Materials list when developing small-scale or microscale procedures for these tests. Note that a total of 2 g of each solid and 20 mL of each solvent are provided for testing.
4. To test the melting point of a substance, first place a small amount of each solid in separate locations in an aluminum evaporating dish. Hold the dish above a boiling water bath with tongs and observe if the solids melt at < 100 °C.
5. If the substance did not melt at < 100 °C, place a pea-size amount of solid in a borosilicate glass test tube. Heat the test tube in a medium burner flame for 1–2 minutes. Record observations. To test the melting point in a medium burner flame, start with a fuel-rich yellow flame and adjust the air inlet until all yellow just disappears. The result is a light blue flame without an inner cone; the temperature is < 500 °C.
6. Record the results of qualitative testing in a data table.
   

Guided-Inquiry Design and Procedure

Development of a Qualitative Analysis Scheme
A nearby school has discovered a set of chemicals that are missing labels. The science teacher has recovered the potential missing labels, but needs to match them with the correct bottles. Your lab has been asked to design a procedure to identify the 12 unknown chemicals. You may use any series of tests that deal with the properties of solids.

Form a working group with other students and discuss the following questions.

1. Compare your group’s results from the Introductory Activity with those of other groups. From these discussions and your test data, list general physical properties that can be associated with each type of bonding in a solid.
2. Using yes−no logic, create a flow chart that can be used to characterize an unknown solid as ionic, polar covalent, nonpolar covalent or metallic.
3. If given a white solid, what testing results would help you to identify the solid as polar covalent?
4. Each group will be given six unknown solids to evaluate. Using the flow chart as a guide, write a detailed, step-by-step procedure for testing the solids and identifying the type of bonding in each. Include the materials and equipment that will be needed, safety precautions that must be followed, amounts of chemicals to use, etc.
5. Carry out the flow chart tests on the six unknowns and record the results of each test in an appropriate data table.
6. Identify each solid as ionic, polar covalent, nonpolar covalent or metallic.
7. Share your group’s data with a group that analyzed the other six unknowns.
8. Copies of the missing labels have been assembled on a sheet of paper. Compare the results of testing the unknowns with the chemical labels. As a class, match the labels with as many unknowns as possible.
9. Perform any additional tests as needed to verify the identity of any remaining chemicals whose identities are ambiguous.

Student Worksheet PDF

14002_Student1.pdf