Materials Included In Kit

Additional Materials Required

Safety Precautions

Hexane is a flammable organic solvent and a dangerous fire risk. Keep away from flames, heat and other sources of ignition. Cap the solvent bottle and work with hexane in a fume hood or designated work area well away from any Bunsen burners used in the lab. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles and chemical-resistant gloves and apron. Please consult 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

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 hexane solutions should be collected in a flammable organic waste container and allowed to evaporate according to Flinn Suggested Disposal Method #18a. All other solids and solutions may be disposed of in the trash according to Flinn Suggested Disposal Methods #26a and #26b, respectively.

Lab Hints

• The laboratory work for this experiment can easily be completed in a typical 50-minute lab period. The Preab Questions may be assigned as homework in preparation for lab or may be used as part of a class discussion prior to doing the lab.
• 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. Using common household materials removes one (unnecessary) stumbling block in this process.
• Many other common solids may also be used. Any metal may be used instead of aluminum and many different ionic compounds may be substituted for sodium chloride. Suitable nonpolar organic solids that may be used instead of or in addition to stearic acid include lauric acid or paraffin wax.
• If several lab sections will be performing this experiment the same day, keep the boiling water baths (step 1) set up throughout the day. Use distilled or deionized water for best results and replenish the boiling stones as needed.
• Borosilicate–glass test tubes are provided in this kit for use in both step 5, testing the solubility of the solids in hexane, and step 11, testing the melting points of the solids in a Bunsen burner flame. Have students dispose of the hexane from step 6 in a flammable, organic waste container, then clean and dry the test tubes. When performing the melting points of the solids in the Bunsen burner flame, make sure students use borosilicate test tubes.
• 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 LabPro or CBL-2 computer interface system.
• Remind students not to use flammable organic solvents around or near a heat source.
• Using a conventional 110-V “lightbulb-type” conductivity tester will require larger sample sizes. It is recommended that the teacher perform the conductivity tests as a demonstration if 110-V conductivity testers will be used.
• See the Supplementary Information in the Further Extensions section for a description of the Mohs hardness scale. (The information may be used as an optional student handout, if desired.) The Mohs hardness scale is a nonlinear, semiquantitative tool that is used in geology to rate the relative hardness of rocks and minerals. The scale ranges from 1 (talc) to 10 (diamond)—the higher the number, the harder the material. An object will only scratch something with a lower hardness rating.
• A lab station can be set up for testing the hardness of various minerals, such as rock candy (sucrose), rock salt (sodium chloride), a candle (paraffin wax) and an aluminum strip. Test the hardness of each solid by trying to scratch them with a fingernail, a penny and an iron nail. Have students record their observations and relative order of hardness of the solids.
• The following demonstration provides a good discrepant event to describe the hardness test. Ask students to predict what will happen if a nail is scraped across the glass stage on the overhead projector. After students have given their dire predictions, rub a nail back and forth on the overhead. The nail will not scratch the glass—steel (iron) has a hardness of 5 while glass has a hardness of 6. Always test this demonstration in one corner of the overhead projector first, however.
• Caution students to use the proper technique to detect the odor of a substance. Place the open container about 6 inches away from the nose and use your hand to waft the vapors toward the nose. While the chemicals in step 3 post no risk from inhalation, every opportunity should be taken to develop safe laboratory techniques.

Teacher Tips

• 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.
• It is hard to convey the principles of bonding and structure using only two-dimensional drawings or pictures. We strongly encourage the use of three-dimensional models to help students recognize and understand the relationship between structure and bonding. Consult your current Flinn Scientific Catalog for a complete selection of models, including diamond (AP6176), graphite (AP6175), ice (AP6178) and sodium chloride (AP6179).
• Many new terms and definitions are introduced in this activity, which provides an overview of all types of chemical bonding. Encourage students to make a list of all the new terminology and write out their definitions. Remind students also to consult their textbooks for additional examples, models and illustrations that may help explain the concepts.

Further Extensions

Supplementary Information: Mohs Hardness Scale

Hardness is not an intrinsic or fundamental physical property of a substance. It is a defined property which can only be assessed by comparing the relative properties of two or more substances. Hardness is useful in mineralogy for the field identification of rocks and minerals.

Hardness is defined as the resistance of a mineral to being scratched. (This is different than breaking or shattering a mineral.) The geologist Friedrich Mohs developed a convenient scale for ranking minerals with respect to hardness. The principle behind the scale is quite simple—an object will only scratch something with a lower hardness rating. The scale and some common comparison tools are listed in the Table 1. Despite the obvious simplicity of the method, the scale actually gives pretty specific results. Thus, a penny will scratch a halite (salt) crystal while a fingernail will not.

Table 1. Mohs Hardness Scale

Hardness testing is extremely important in materials science and engineering for steel and other alloys, ceramics and even plastics. Modern methods, such as the Rockwell hardness test, measure the depth or area of an indentation left by a diamond cone or a steel ball when a measured force is applied for a specified period of time.

Answers to Prelab Questions

1. A student wanted to illustrate the structure of magnesium chloride and decided simply to replace the Na⁺ ions in Figure 1 (Background section) with Mg²⁺ ions. What would be wrong with the resulting picture?

   The picture would show the wrong ratio of ions in the crystal structure. The formula of magnesium chloride is MgCl₂—there are two chloride ions for every magnesium ion. The ratio of positive and negative ions in the sodium chloride crystal structure is 1:1.

2. 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?

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

3. The three dimensional structure of diamond, a crystalline form of the element carbon, is shown in Figure 3. Use this structure to explain why diamond is the hardest known material.
   {13964_PreLab_Figure_3}

   Diamond is a covalent-network solid. The structure consists of strong covalent carbon–carbon single bonds in all directions. Each carbon atom forms four bonds and thus has a stable octet of valence electrons. Cutting a diamond would require breaking many carbon–carbon bonds.

Sample Data

*The average temperature of a Bunsen burner flame is greater than 1000 °C. Microburners may not have as high a flame temperature.
†The melting point of sodium chloride (801 °C) is greater than that of pure aluminum metal (660 °C). Sodium chloride is observed to melt in a test tube placed in a Bunsen burner flame, while aluminum granules generally do not melt under these conditions. This is probably due to the invisible oxide coating which is always present on aluminum. The melting point of aluminum oxide is about 2000 °C.

Answers to Questions

1. Compare the volatility and odor of stearic acid and sucrose. Which is more volatile? Why? Is it possible for a compound to be volatile but have no odor? Explain.

   Stearic acid has an odor and seems to be more volatile than sucrose. In order for a substance to have an odor, some molecules must enter the gas phase and diffuse in air to reach the nose. Some volatile substances, however, may not have an odor, because the nose lacks the appropriate receptors to “detect” the odor.

2. Both stearic acid and sucrose are molecular substances, but one is polar and the other is nonpolar. Compare the solubility of the two compounds in water and in hexane to determine which is which.

   Stearic acid dissolved in hexane, not in water. Sucrose dissolved in water, not in hexane. This suggests that stearic acid is nonpolar (like hexane), while sucrose is polar (like water). Note to teachers: Stearic acid consists of a very long (C₁₇H₃₄—), nonpolar hydrocarbon “tail” attached to a small polar carboxylic acid (—CO₂H) group. The nonpolar hydrocarbon tail dominates the physical properties of the solid (e.g., solubility, melting point).

3. Based on the answers to Questions 1 and 2, predict whether the intermolecular forces (forces between molecules) are stronger in polar or nonpolar substances.

   Polar substances have stronger intermolecular forces—it takes more energy to pull polar molecules apart and have molecules enter the gas phase.

4. In order for a substance to conduct electricity, it must have free-moving charged particles.
   1. Explain the conductivity results observed for sodium chloride 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.

5. Complete the following table:
   {13964_Answers_Table_1}
   Note: Stress to students that these are general properties—there are many exceptions. The melting points of metals, for example, range from –39 °C (for mercury) to 3407 °C (for tungsten). Many low-melting metals (lithium, sodium, potassium, gallium, etc.) are also soft enough that they can be cut with a knife. Finally, not all ionic compounds are water-soluble.

References

This experiment has been adapted from Flinn ChemTopic™ Labs, Volume 5, Chemical Bonding; Cesa, I., Ed., Flinn Scientific: Batavia, IL, 2003.

Introduction

Looking for patterns in the properties of different substances can help us understand how and why atoms join together to form compounds. What kinds of forces hold atoms together? How does the nature of the forces holding atoms together influence the properties of a material?

Concepts

• Chemical bonds
• Ionic bonding
• Covalent bonding
• Metallic bonding

Background

Groups of atoms are held together by attractive forces that are called 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 is used to describe the 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 arrange themselves 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 equally 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, which must exist therefore 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 experiment is to study the physical properties of common solids and to investigate the relationship between the type of bonding in a substance and its properties. The following physical properties will be studied:

• Volatility and Odor: Volatile substances evaporate easily and may have an odor.
• Melting Point: The temperature at which a solid turns into a liquid.
• Solubility: Ability of one substance to dissolve in another. Water is a highly polar solvent. Hexane is a nonpolar solvent.
• Conductivity: Ability to conduct electricity.
• Brittleness: Tendency of a solid to break or crumble when a stress is applied.

Materials

Prelab Questions

1. A student wanted to illustrate the structure of magnesium chloride and decided simply to replace the Na⁺ ions in Figure 1 (Background section) with Mg²⁺ ions. What would be wrong with the resulting picture?
2. 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?
   2. Which is more polar, an O—H or N—H bond?
3. The three-dimensional structure of diamond, a crystalline form of the element carbon, is shown in Figure 3. Use this structure to explain why diamond is the hardest known material.
   {13964_PreLab_Figure_3}

Safety Precautions

Hexane is a flammable organic solvent and a dangerous fire risk. Keep away from flames, heat and other sources of ignition. Cap the solvent bottle and work with hexane in a fume hood or designated work area well away from the Bunsen burner used. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the lab.

Procedure

1. Prepare a boiling water bath for use in step 10: Half-fill a 150-mL beaker with water, add a boiling stone and heat the beaker on a hot plate or Bunsen burner setup at a medium setting.
2. Label five weighing dishes for the five solid samples and obtain 0.2–0.3 g samples of each solid in the appropriate weighing dish. Record the color and appearance of each solid in the data table. The five samples are: aluminum, silicon dioxide, sodium chloride, stearic acid, and sucrose.
3. Test the volatility and odor of each solid by wafting any vapors to your nose with your hand. Record all observations in the data table. Note: To detect the odor of a substance, place the open container about 6 inches away from the nose and use your hand to waft the vapors toward the nose.
4. Test the conductivity of each solid by touching the wires of the conductivity tester directly to the solid. Record the conductivity of each sample in the data table.
5. Place five small test tubes in the test tube rack. Label the five small test tubes for the five solid samples and add a small amount of each solid, about the size of a grain of rice, to its labeled test tube.
6. Add about 20 drops of hexane to each test tube. Stir each mixture and observe whether the solid dissolves in hexane. Record the results in the data table. Dispose of the hexane as directed by your instructor, then clean and dry the test tubes.
7. Add about 20 drops of water to each weighing dish. Stir each mixture and observe whether the solid dissolves in water. Record the solubility (soluble, partially soluble or insoluble) in the data table.
8. For water-soluble substances only: Determine the conductivity of the aqueous solution by placing the wires of the conductivity tester directly into the liquid. Record the results in the data table.
9. Obtain a weighing dish and place a small, pea-sized amount of each solid in separate locations on the dish.
10. Using the test tube clamp, set the dish on top of the boiling water bath and heat the solids for 1–2 minutes. Observe whether any of the solids melt and record the observations in the data table.
11. For solids that did not melt at the boiling water bath temperature: Place a small, pea-sized amount of each solid in a clean and dry test tube. Using a test tube holder, heat the test tube in a burner flame for 1–2 minutes. Record observations in the data table.
12. Test the brittleness of each solid by placing a small sample in the mortar designated for it and grinding with the pestle. Record the observations in the data table.
13. Dispose of all waste liquids and solids as directed by your instructor.

Student Worksheet PDF

13964_Student1.pdf