Materials Included In Kit

Additional Materials Required

Safety Precautions

Use only borosilicate (e.g., Pyrex^®) test tubes for this laboratory. Inspect all test tubes and do not use any cracked or chipped test tubes. Warn students to work carefully to avoid scalding themselves with the hot water bath. Cetyl alcohol is slightly toxic by ingestion. 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 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 organic solids may be disposed of according to Flinn Suggested Disposal Methods #18b or #24a. Alternatively, the solids may be recycled from class to class and also from year to year.

Lab Hints

• The laboratory work for this experiment can easily be completed in a typical 50-minute class period. If desired, however, the lab may be simplified by dividing the class into groups of four students, with two “working pairs” in each group. Each working pair would be responsible for only one part of the experiment, either Part A or Part B. The two working pairs would then share the data with each other to complete the Post-Lab Questions. Two groups may share one hot plate or hot water bath.
• Excellent cooling curve data is obtained in Part A with as little as 2–3 g of solid in a medium test tube. We recommend using digital thermometers for best results. In order to get good results in Part B, however, a larger sample size is recommended. When smaller sample sizes were tested in Part B, the temperature of the solid steadily increased when placed in the hot water bath and the heating curve did not exhibit a well-defined “plateau” at the melting point. Three variables—the amount of solid, the temperature of the hot water bath and the amount of water in the hot water bath—will influence the quality of the heating curve data.

Teacher Tips

• “Phase-change wallboard” has been developed as a passive thermal storage construction material to improve energy efficiency in heating and cooling buildings. Phase-change wallboard contains paraffin wax embedded in gypsum. The paraffin undergoes reversible phase changes when heated or cooled, absorbing or releasing large amounts of heat and maintaining a constant temperature in the process. Consider how a phase-change material (PCM) might be used in buildings in a mild-climate area. During the day, when the outside temperature increases, the solid PCM melts and absorbs (stores) the excess heat energy. This cools the building and reduces the need for external cooling (air conditioning). The reverse occurs at night—when the outside temperature decreases, the liquid PCM solidifies and releases its “stored” heat energy.

Answers to Prelab Questions

1. The kinetic-molecular theory (KMT) describes how close together the molecules are in a solid, liquid and a gas; their relative motion; and the attractive forces between the molecules. Use the KMT to explain the following properties of liquids and solids:
   1. A liquid flows and takes the shape of its container.

      The molecules in a liquid remain close together, but move freely and randomly.

   2. Solids are generally incompressible.

      The molecules in a solid are physically as close together as possible (tightly packed).

   3. Liquids have a definite volume.

      The molecules in a liquid are close together and there are attractive forces between the molecules.

   4. A solid absorbs heat from its surroundings as it melts.

      Relatively strong attractive forces exist between the molecules in the solid state—these forces must be (partially) broken when the solid changes to a liquid.

2. The following graph shows heating curve data for ice, water, and steam as heat energy is added to the system at a constant rate. (a) In what regions of the curve (A–E) is the average kinetic energy of the molecules increasing? (b) In what region of the curve are ice and water present together? (c) What happens to the heat energy that is absorbed by the molecules in this region of the curve?
   {12812_Answers_Figure_2}
   1. The average kinetic energy of the molecules increases as the temperature increases in regions A, C and E.
   2. Ice and water are present together at the melting point (freezing point) of water (0 °C, region B).
   3. The heat energy absorbed by the ice molecules is used to break the attractive forces between the molecules in the crystal lattice. Note: The average kinetic energy of the molecules remains constant. The heat energy increases the potential energy of the molecules. At the same temperature, molecules in the liquid state have a higher potential energy than molecules in the solid state.

Sample Data

* The cooling and heating curve trials may be stopped when the temperature reaches 30 °C and 55 °C, respectively.

Answers to Questions

1. Prepare a graph of temperature on the y-axis versus time on the x-axis. Plot the data from Parts A and B as two series of points, using different color pencils or different shapes to mark the points for Part A versus Part B. Draw a smooth (continuous) curve through the plotted points for each series of data.
   {12812_Answers_Figure_3}
2. Label the following regions (A–C) on the cooling curve: A, only liquid is present; B, liquid and solid are present together; C, only solid is present.
3. What happens to the temperature of a pure substance while it is freezing or melting? Estimate the freezing point and the melting point of the organic solid from the cooling curve and the heating curve, respectively. Does the freezing point/melting point depend on the direction in which the phase change takes place?

   The temperature of a pure substance should remain constant at the freezing point or melting point as long as both phases are present. The freezing point of the melted solid is estimated from the “flat” region of the cooling curve—44.0 °C. The heating curve did not “level off” at a particular temperature as did the cooling curve. However, the rate at which the temperature increased slowed down noticeably in the 43–44 °C temperature range, giving an estimate of 43.5 °C for the melting point. The slight difference in the estimated freezing point/melting point is not significant (within experimental error). The freezing point/melting point does not depend on the direction in which the phase change takes place.

4. Use the following information to identify the unknown organic solid.

   Cetyl alcohol, C₁₆H₃₃OH, mp 54–56 °C
   Lauric acid, C₁₁H₂₃CO₂H, mp 43–44 °C
   Stearic acid, C₁₅H₃₃CO₂H, mp 67–69°C
   The unknown solid had a melting point of 43–44 °C—lauric acid.

5. Which set of data (the cooling curve or the heating curve) provided a more accurate or a more precise estimate of the melting point? Which variables in the design of the experiment might account for any difference in the results?

   The literature melting point of lauric acid is 43–44 °C. The heating curve data gave a more accurate estimate of the melting point, but the cooling curve data gave a more precise value. The cooling curve was qualitatively easier to interpret because there was a true temperature plateau at the freezing point. The amount of lauric acid was the same in both parts of the experiment. Several variables, however, were different in Part A versus Part B. The hot water bath contained more water than the cold water bath and was therefore a greater heat “reservoir.” The solid sample in Part B could not be stirred, making heat transfer less efficient. The cold water bath was insulated in a coffee cup, the hot water bath was not. Note: This is a great discussion question—students usually grasp at straws when it comes to analyzing experimental error. Ask students to consider how these variables would affect the rate at which heat is transferred into or out of the system, and how the temperature versus time data would be affected as a result.

6. Circle the correct choices: Freezing is an exothermic process—the liquid releases heat to its surroundings. At the freezing point, the average potential energy of the molecules decreases and the liquid solidifies.

   Note: Question 6 may be difficult for students to grasp. Energy is released during freezing but the temperature does not change. Average kinetic energy is dependent on temperature so it does not change. The potential energy must decrease and this is due to the formation of intermolecular forces holding the solid together. Molecules in the liquid state have a higher potential energy than molecules in the solid state.

7. Answer increases, decreases or no change to predict how doubling the amount of solid would change the results in Part B:
   1. The rate at which the temperature of the solid increases. ___Decreases___
   2. The temperature at which the solid melts. ___No change___
   3. The amount of heat absorbed by the sample as it melts. ___Increases___
8. (Optional) Lauric and stearic acid are fatty acids (components of fats and oils) that are used to make soap. What factor might explain the regular increase in the melting point of the following fatty acids as the number of carbon atoms increases?

   Lauric acid, CH₃(CH₂)₁₀CO₂H, mp 43.2 °C
   Myristic acid, CH₃(CH₂)₁₂CO₂H, mp 54.0 °C
   Palmitic acid, CH₃(CH₂)₁₄CO₂H, mp 61.8 °C
   Stearic acid, CH₃(CH₂)₁₆CO₂H, mp 68.8 °C
   In general, the greater the attractive forces between the molecules in a molecular solid, the higher its melting point. This implies that the strength of the attractive forces increases as the number of carbon atoms increases. Note: The properties of the fatty acids are dominated by the long nonpolar hydrocarbon “tail.” The principal attractive forces acting between nonpolar molecules are London dispersion forces. The strength of London dispersion forces increases as the size of the molecules increases.

References

This kit was adapted from Solids and Liquids, Flinn ChemTopic™ Labs, Vol. 11, Cesa, I., Editor; Flinn Scientific Inc.: Batavia, IL (2005).

Introduction

When freezing weather is predicted, Florida’s orange growers spray their trees with water to prevent the fruit from freezing. As the water freezes, it releases heat to the surroundings and protects the fruit from damage. The temperature of the freezing water mixture will remain at the freezing point as long as both ice and water are present. Let’s look at the temperature changes and the energy changes that take place when a liquid freezes or a solid melts.

Concepts

• Solids and liquids
• Phase changes
• Melting point
• Kinetic-molecular theory

Background

The temperature changes and energy changes that occur when a solid melts or a liquid freezes can best be understood by imagining what solids and liquids look like at the level of molecules or ions. Solids and liquids differ in how ordered or rigid their structures are and in the range of motion that the molecules or ions are allowed. Molecules in a crystalline solid are packed together in an ordered three-dimensional pattern, called a crystal lattice, where they are “held in place” by attractive forces between the molecules. The motion of molecules in the solid state is limited to vibrations (stretching and rocking motions)—the molecules are not free to move away from their fixed positions. The forces between molecules in the liquid state are not well understood. Molecules in the liquid state are free to move and are not locked in position. Attractive forces between molecules, however, tend to keep the molecules close together, so that their motion is perhaps best described as coordinated rather than independent.

A solid and its liquid are in equilibrium at the melting point, the temperature at which a crystalline solid becomes a liquid. The melting point of a pure substance is a characteristic physical property that can be used to identify a substance and to determine its purity. When a solid is heated, the temperature of the solid will increase until it reaches the melting point. Temperature is related to the average kinetic energy of the molecules—as the temperature increases, the average kinetic energy of the molecules increases and they begin to vibrate more rapidly. At the melting point, the vibrations become so rapid that the molecules begin to “break loose” from their fixed positions and melting occurs. The temperature of the solid–liquid mixture will remain constant at the melting point until all the solid has melted. Although the temperature remains constant at the melting point, heat energy must be added to break the attractive forces between molecules. In general, the more orderly the packing arrangement of molecules in the solid state and the stronger the attractive forces between molecules, the higher the melting point will be and the more heat that will be needed to melt the solid. The reverse process occurs when a liquid freezes. When a liquid freezes, energy in the form of heat is released to the surroundings. The same amount of heat required to melt a solid will be released by the liquid when it freezes or solidifies.

Experiment Overview

The purpose of this experiment is to investigate the solid–liquid phase changes for an unknown organic solid. Temperature versus time data will be measured as the melted solid is slowly cooled (Part A), and again as the fused solid is reheated (Part B). The data will be graphed and the heating and cooling curves will be analyzed to determine the melting point of the solid and identify the unknown.

Materials

Prelab Questions

1. The kinetic-molecular theory (KMT) describes how close together the molecules are in a solid, liquid, and gas, their relative motion, and the attractive forces between the molecules. Use the KMT to explain the following properties of liquids and solids:
   1. A liquid flows and takes the shape of its container.
   2. Solids are generally incompressible.
   3. Liquids have a definite volume.
   4. A solid absorbs heat from its surroundings as it melts.
2. The following graph shows heating curve data for ice, water and steam as heat energy is added to the system at a constant rate. (a) In what regions of the curve (A–E) is the average kinetic energy of the molecules increasing? (b) In what region of the curve are ice and water present together? (c) What happens to the heat energy that is absorbed by the molecules in the region where ice and water are both present?
   {12812_PreLab_Figure_1}

Safety Precautions

Read the entire Procedure before beginning the experiment. Work carefully to avoid scalding yourself with hot water. Cetyl alcohol is slightly toxic by ingestion. 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 laboratory.

Procedure

Part A. Cooling Curve

1. Fill a 400-mL beaker two-thirds full with hot tap water. Heat the water to about 80 °C on a hot plate. Proceed to steps 2–5 as the water is heating.
2. Obtain about 6 g of the “unknown solid” in a weighing dish and transfer the solid to a clean and dry test tube.
3. Add about 100 mL of cold tap water (15–20 °C) to a Styrofoam cup and nest the cup inside a second Styrofoam cup. Place the nested Styrofoam cups in a 250-mL beaker and set the beaker on the ring stand.
4. Hold the test tube (prepared in step 2) with a clamp or test tube holder and place the test tube into the hot water bath (step 1) at 80 °C.
5. Insert a digital thermometer into the solid in the test tube. When the temperature of the solid is about 80 °C, remove the test tube from the hot water bath and clamp the test tube to the ring stand.
6. Measure and record the precise temperature of the melted organic solid in the test tube and then immediately lower the test tube into the cold water bath in the Styrofoam cup. Start timing.
7. Carefully stir the melted solid with the digital thermometer and measure the temperature every 30 seconds for 10 minutes, or until the temperature is about 30 °C (whichever comes first). Record all time and temperature readings in the data table. Note: Continue stirring the sample until it is no longer possible to do so (when the sample has solidified).

8. Check that the temperature of the hot water bath is between 75 and 80 °C. Do not let the temperature rise above 80 °C—add cold water if necessary to adjust the temperature.
9. Remove the test tube from the cold water bath (step 7). Measure and record the precise temperature of the solid and immediately place the test tube in the 80 °C water bath. Start timing.
10. Measure and record the temperature of the unknown organic solid every 30 seconds for 10 minutes, or until the temperature begins to rise steeply (whichever comes first). Note: Start stirring once the sample has softened up enough to do so, but do not try to “force” it.
11. Remove the test tube from the hot water bath and dispose of the sample as directed by the instructor.

Student Worksheet PDF

12812_Student1.pdf