Materials Included In Kit

Additional Materials Required

Safety Precautions

Hexane is a highly flammable liquid and vapor and a dangerous fire risk. Keep away from flames, heat and other sources of ignition. Hexane is an aspiration hazard—it may be fatal if swallowed and enters airways. It also causes skin and eye irritation and may cause drowsiness or dizziness if inhaled. Hexane is a suspected reproductive toxin. Do not use if you are pregnant. Cap the solvent bottle and work with hexane in a fume hood; make sure it is not used near or in the same area as the Bunsen burner used in step 13. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information. Remind students to wash their hands thoroughly with soap and water before leaving the lab.

Disposal

Consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The hexane solutions should be collected in a designated flammable organic waste container for licensed hazardous waste disposal according to Flinn Suggested Disposal Method #18b. All other solids and solutions may be disposed of in the trash or drain according to Flinn Suggested Disposal Methods #26a and b, respectively.

Lab Hints

• The laboratory work for this experiment can easily be completed in a typical 2-hour lab period. The Prelaboratory Assignment should be reviewed as part of a class discussion prior to 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 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 and paraffin wax.
• The activity may be extended by increasing the complexity and number of solids. Add a polar covalent solid, such as glycine or dodecyl alcohol (the latter is very low-melting), and incorporate a metal oxide, such as iron oxide, as well as a semiconductor (e.g., silicon or even graphite).
• 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-V battery, the conductivity tester is convenient, portable, and safe. Conductivity tests may also be done using conductivity sensors.
• Using a conventional 110-V “lightbulb-type” conductivity tester will require larger sample sizes. It is recommended that the instructor perform conductivity tests as a demonstration if a 110-V tester 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.
• 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 an overhead projector or document camera. After students have given their dire predictions, rub a nail on the glass surface. 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 a small inconspicuous corner of the stage before doing it for a class!
• Individual mineral samples (halite and quartz) suitable for hardness testing are available from Flinn Scientific.

Teacher Tips

• Students may share balances, conductivity testers, hot plates (boiling water baths), laboratory burners and mortars and pestles. Set up five mortars in a designated location and label them with the names of the solids to be tested.

Further Extensions

Supplementary Information

Mohs Hardness Scale
Hardness is not an intrinsic or fundamental physical property of a substance. It is a defined property that 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 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.

Hardness testing is 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 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 (2.5 and 2.2, 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 (3.5–2.2) than that between N and H (3.0–2.2).
      An O—H bond is more polar than an N—H bond.

3. The three dimensional structure of diamond, a beautifully crystalline form of the element carbon, is shown in Figure 3. Use this structure to explain why diamond is the hardest known material. What kind of solid is this?
   {14030_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.
4. In order for a substance to conduct electricity, it must have free-moving charged particles.
   1. Use the model of metallic bonding to explain why metals conduct electricity.

      Metals conduct electricity because the valence electrons are delocalized among the metal cations in the crystal structure and are able to move freely throughout the solid.

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

      Molten sodium chloride is expected to conduct electricity because the particles (ions) in a liquid are in constant, random motion.

Sample Data

Laboratory Report

*The average temperature of a Bunsen burner flame is greater than 1000 °C. Microburners will have a lower 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 due to the invisible oxide coating on aluminum. The melting point of aluminum oxide is about 2000 °C.

Mineral Hardness**
(Optional) Use this space to record observations of the hardness of mineral samples.

A candle can be scratched with a fingernail (Mohs hardness = 2).
Rock candy (sucrose) may also be scratched with a fingernail (Mohs hardness = 2).
Aluminum metal and a salt crystal can be scratched with a penny but not with a fingernail (Mohs hardness = 4).
A quartz (silicon dioxide) crystal can be scratched with a ceramic pestle but not with an iron nail (Mohs hardness = 7).

** See the Supplementary Information for a description of hardness testing and the Mohs hardness scale.

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: Stearic acid consists of a long, nonpolar hydrocarbon (C₁₇H₃₅—) “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 apart” polar molecules and have molecules enter the gas phase.

4. 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 conducts electricity because the ions are no longer fixed into position. (The solute particles in a liquid are able to move freely.)

5. Name the three hardest substances that were tested. To what classes of solids do these substances belong? What general feature do these three types of solids have in common?

   The three hardest substances were aluminum (a metal), silicon dioxide (a covalent-network solid), and sodium chloride (an ionic compound). All of these solids have extended (infinite), three-dimensional crystal structures with strong bonds in all directions. There are no individual molecules in the solid state.

6. Compare the hardness and brittleness of aluminum versus salt. Suggest a reason, based on the crystal structure of metals versus ionic compounds, why hardness and brittleness are not the same thing.

   Aluminum is hard and nonbrittle. Salt is hard and brittle. The hardness of both solids is probably due to their extended, threedimensional crystal structures. There are strong bonds in all directions and it is hard to apply enough force to break the bonds and dislocate an atom (or ion). Brittleness relates to what happens when particles in the solid state have been dislocated. The oppositely charged ions in an ionic lattice must occupy specific positions for optimum ionic bond strength. If some ions are displaced by force, attractive forces may be replaced by repulsive forces between ions of like charge. As a result, an ionic solid breaks or crumbles easily. In a metal, however, it does not matter if metal atoms are displaced by force because all of the atoms are identical and the electrons can move around to minimize repulsion of the metal cations.

7. Complete the following table summarizing the general properties of various solids.
   {14030_Answers_Table_2}
   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 (e.g., lithium, sodium, potassium, gallium) are also soft enough that they can be cut with a knife. Finally, not all ionic compounds are water soluble.
8. Fill in the blanks in the sample flow chart to show how different solids may be identified using a sequence of simple laboratory tests.
   {14030_Answers_Figure_4}

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
• Polar bonds
• Conductivity

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 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 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. 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 consist of atoms forming covalent bonds with each other in all directions. The result is an almost infinite network of strong covalent bonds—with 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 two electrons 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 depicted 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 nonpolar.
• Conductivity: Ability to conduct electricity.
• Hardness: Resistance of a substance to being scratched.
• 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 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 beautifully crystalline form of the element carbon, is shown in Figure 3. Use this structure to explain why diamond is the hardest known material. What kind of solid is this?
   {14030_PreLab_Figure_3}
4. In order for a substance to conduct electricity, it must have free-moving charged particles.
   1. Use the model of metallic bonding to explain why metals conduct electricity.
   2. Would you expect molten sodium chloride to conduct electricity? Why or why not?

Safety Precautions

Hexane is a highly flammable liquid and vapor and a dangerous fire risk. Keep away from flames, heat and other sources of ignition. Hexane is an aspiration hazard—it may be fatal if swallowed and enters airways. It also causes skin and eye irritation and may cause drowsiness or dizziness if inhaled. Hexane is a suspected reproductive toxin. Do not use if you are pregnant. Cap the solvent bottle and work with hexane in a fume hood; make sure it is not used near or in the same area as the Bunsen burner used in step 13. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the lab.

Procedure

1. Prepare a boiling water bath for use in step 12: Half-fill a 150-mL beaker with water, add a boiling stone, and heat the beaker on a hot plate at a medium setting.
2. Label five weighing dishes 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.
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.
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. Obtain a 24-well reaction plate and add a small amount of each solid (about the size of a grain of rice) to separate wells A1–A5, in the order shown in the data table.
6. Add about 20 drops of water to each well. Stir each mixture and observe whether the solid dissolves in water. Record the solubility using the following terms—soluble, partially soluble or insoluble.
7. For water-soluble substances only: Determine the conductivity of the aqueous solution by placing the wires directly into the liquid. Record the results in the data table.
8. Label five small test tubes or a ceramic spot plate and add a small amount of each solid, about the size of a grain of rice, to separate test tubes. Place the test tubes in a rack.
9. Take the test tube rack to the hood and add about 20 drops of hexane to each tube. Stir each mixture and observe whether the solid dissolves in hexane. Record the observations and results.
10. Discard the contents of the test tubes into a designated flammable waste container as directed by the instructor. Rinse the test tubes with acetone or another solvent.
11. Obtain a large, disposable aluminum evaporating dish or Pyrex^® watch glass and place a small, pea-sized amount of each solid in separate locations on the dish.
12. Set the dish on top of the boiling water bath and heat the dish with the solids for 1–2 minutes. Observe whether any of the solids melt and record observations in the data table.
13. For solids that did not melt at the boiling water bath temperature: Place a small, pea-size amount of each solid in a clean, dry Pyrex 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.
14. Test the brittleness of each solid by placing a small sample in the mortar designated for it. Grind the solid with the pestle. Record observations in the data table.
15. (Optional) Test the hardness of mineral samples by trying to scratch them with a fingernail, penny and nail. Record observations. Note: Your instructor may distribute a supplementary handout on hardness testing.

Student Worksheet PDF

14030_Student1.pdf