Materials Included In Kit

Additional Materials Required

Safety Precautions

Potassium nitrate in solid form is a strong oxidant and a fire and explosion risk when heated or in contact with organic materials. It is also a skin irritant. Avoid contact with skin and eyes. Wear chemical splash goggles and chemical-resistant gloves and apron. Use caution when working with a hot water bath and a hot plate. Wear heat-resistant gloves or a heat protector when handling hot glassware. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory. Please consult 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. Potassium nitrate solutions may be rinsed down the drain with excess water according to Flinn Suggested Disposal Method #26b. Alternatively, the solutions may be collected in a central location and allowed to evaporate. The solid potassium nitrate remaining after solvent evaporation may be recrystalized for use by another class or for the following year.

Lab Hints

• The laboratory work for this experiment can reasonably be completed in one 50-minute lab period. The Prelab Questions may be assigned as homework in preparation for lab or they may be used as the basis of a cooperative class discussion before lab.
• Adjust the working group size and the number of samples each group tests to accommodate your lab setting and schedule. The experiment as written calls for each pair of students to test three solutions and share their data with another pair of students in the same working group. Doing the experiment this way gives six data points for the solubility curve. The solubility trend is apparent with as few as four data points, but the best-fit curved line will give better agreement with the literature if more points are used. Teachers should grade more by the shape of the solubility curve than by the solubility values. Alternatively, the experiment may be done as a cooperative activity by plotting class data. This approach may give the best overall fit for the solubility curve.
• Have students use warm tap water in their water baths. This will help speed up the heating process and save time. Remind students not to boil the water in their hot water bath. Temperatures of 80–90 °C work well.
• If ice is available in the lab, students may extend the solubility curve downward below room temperature by decreasing the mass of potassium nitrate. Use mass ratios of 0.20–0.30 g KNO₃ per gram of water to obtain saturation temperatures between 0 and 20 °C. Adding a couple of data points in this temperature range will improve the fit between the literature and the experimental solubility curve.
• Temperature measurements may be made using digital thermometers, glass-bulb thermometers, or computer-interfaced temperature sensors. Digital thermometers are preferred over glass thermometers because they provide direct readings, update every second, and have a precision of ±0.1 °C. Glass thermometers are fragile and easily broken, especially if the solutions are vigourously stirred, as suggested in the Sample Procedure. In addition, the 1 °C divisions that are marked on most glass thermometers make them less precise (±0.5 °C) than digital thermometers. Never allow students to use a glass thermometer as a stirring rod.
• Distribute magnifying glasses, if available, to help students observe the first signs of crystal formation. Test tubes must be clean for best results. We supply and recommend using reagent rather than laboratory grade potassium nitrate for this experiment. The potassium nitrate may be recovered from the solutions and recrystalized for use in subsequent years.
• Microscale technique is important in steps 6 and 7. The mass of water must be in the 0.9 g–1.1 g range to avoid having too concentrated or too dilute solutions.
• If hot plates are not available, a Bunsen burner setup or immersible water heater can be used.
• Expect quite a bit of scatter in the saturation temperature measurements obtained by different student groups. The Supplementary Information in the Further Extensions section has a sample graph that shows both literature data and typical classroom data for the solubility of potassium nitrate.

Teacher Tips

• The temperature dependence of solubility depends on whether the heat of solution is exothermic or endothermic. For many (although not all) ionic compounds, the heat of solution is endothermic and their solubility increases as the temperature increases. The amount of heat absorbed or released when an ionic compound dissolves in water determines the effect of temperature on the solubility of the salt. Whether heat is absorbed or released depends on both the energy required to disrupt the crystal structure of the ionic solid and the energy liberated when the cations and anions interact with the solvent.
• This experiment may be referred to again when equilibrium is introduced—the temperature dependence of solubility may be explained in terms of Le Chatelier’s Principle and the effect of temperature on the equilibrium constant for the reaction (the solubility product constant).
• A common student misconception is that a saturated solution is at equilibrium because no more solute will go into solution. This is not true—the solution is at equilibrium because the rate at which the solid dissolves is exactly equal to the rate at which the solid crystalizes out.
• Teachers who have access to computer-based graphing programs (such as Vernier’s Graphical Analysis Software, Flinn Catalog No. TC1404) may want to schedule additional lab time for students to graph their data. Students may notice (and point out to their teachers) that they are plotting the independent variable on the y-axis. This is opposite of what they are normally instructed and expected to do.

Further Extensions

Supplementary Information

Answers to Prelab Questions

1. Many solutes, including potassium nitrate, have a tendency to remain in solution even after they have been cooled to below the saturation point. This phenomenon is known as supersaturation. Read the entire Procedure carefully. What measure is taken to prevent supersaturation during this experiment?

   Supersaturation occurs when there are no sites for crystal formation in a hot, unstirred solution. Stir the solution continuously with the thermometer to provide agitation and stimulate crystal formation. Note: Using old test tubes with a scratch or two will also prevent supersaturation. Do not stir using a regular, glass-bulb thermometer.

2. A mixture containing 2.75 g of ammonium chloride (NH₄Cl) in 5.0 g of water was heated to dissolve the solid and then allowed to cool in air. At 61 °C, the first crystals appeared in solution. What is the solubility of ammonium chloride (in g of NH₄Cl per 100 g of water) at 61 °C?

   The ratio of the mass of solute to the mass of solvent is (2.75 g/5.0 g) or 0.55 g of NH₄Cl per gram of H₂O. Multiply this ratio by 100 to obtain the solubility in the desired units of grams of solute per 100 grams of solvent. The solubility of ammonium chloride is 55 g NH₄Cl per 100 g of water at 61 °C.

   {13966_PreLabAnswers_Equation_1}
3. The solubility of ammonium chloride in water was measured as described in this experiment and graphed as follows. Use the solubility curve to predict the solubility of ammonium chloride in water at 40 °C.
   {13966_PreLabAnswers_Figure_1}

   To estimate the solubility of the solute at a specific temperature, draw a straight line from the temperature on the x-axis to the best-fit curved line through the data. Follow this point across horizontally to where it crosses the y-axis. The solubility of ammonium chloride at 40 °C is approximately 44 g per 100 g of water.

4. One of the students doing the experiment only had time to measure the saturation temperatures for three solutions at 51 °C, 61 °C and 80 °C, respectively. Looking at the graph, do you think the student would have been able to accurately predict the shape of the solubility curve based on these three points? Explain.

   Using only three data points, the student might have incorrectly concluded that the relationship between solubility and temperature was linear (followed a straight line). Ideally, the solubility of a substance in water should be measured over its entire liquid range.

Sample Data

Answers to Questions

1. Calculate the mass of potassium nitrate and the mass of water in each solution.

   For sample I-A:

   Mass of KNO₃ = 5.80 – 5.30 = 0.50 g
   Mass of H₂O = 6.77 – 5.80 = 0.97 g

   See the sample results table for the results of other calculations.

2. Calculate the ratio of the mass of potassium nitrate to the mass of water for each solution.

   For sample I-A: Mass ratio KNO₃/H₂O = 0.50 g/0.97 g = 0.52
   See the sample results table for the results of other calculations.

3. Multiply the mass ratio by 100 to determine the concentration of each saturated solution in grams of potassium nitrate per 100 grams of water.

   For sample I-A: 0.52 x 100 = 52 g KNO₃ in 100 g H₂O
   See the sample results table for the results of other calculations.

4. Plot a graph of solubility of potassium nitrate (in g of solute/100 g of water) on the y-axis versus temperature on the x-axis. Scale each axis as necessary. Draw a smooth, best-fit curved line though the data points. Don’t forget to label each axis and give the graph a title!
   {13966_Answers_Figure_3}
5. Using your graph, estimate the solubility of potassium nitrate in water at (a) 0 °C, (b) 50 °C and (c) 100 °C.

   Estimated solubility of potassium nitrate:
   (a) 22 g KNO₃/100 g H₂O at 0 °C
   (b) 85 g KNO₃/100 g H₂O at 50 °C
   (c) >200 g KNO₃/100 g H₂O at 100 °C off-scale on above graph!)

6. Using your graph, predict the temperature at which each of the following mixtures of potassium nitrate in water would form a saturated solution: (a) 15 g KNO₃ in 25 g H₂O; (b) 100 g KNO₃ in 150 g H₂O. Hint: Convert the concentrations to the proper units for solubility before referring to the graph.
   {13966_Answers_Table_4}
7. Define the terms saturated, unsaturated and supersaturated as they apply to solutions. Use complete sentences.

   A saturated solution contains the maximum amount of solute that will dissolve in a given amount of solvent at a particular temperature. A saturated solution usually contains some undissolved solid indicating that no more solid will dissolve at that temperature. An unsaturated solution contains less than the maximum amount of solid that will dissolve in a solvent at a particular temperature. A supersaturated solution contains more than the maximum amount of solute that can dissolve in a given amount of solvent at a particular temperature. Assuming that sufficient time has been allowed to dissolve the solute, neither unsaturated nor supersaturated solutions should contain any undissolved solute.

8. Based on your graph, classify each of the following solutions as either unsaturated or supersaturated at the indicated temperature. Assume that the solutions do not contain any undissolved solid. (a) 50 g KNO₃ in 100 g H₂O at 40 °C; (b) 70 g of KNO₃ in 40 g H₂O at 70 °C. Explain your reasoning.

1. A data point corresponding to a solubility of 50 g KNO₃ per 100 g H₂O and a temperature of 40 °C lies below the best-fit solubility curve drawn through the data points in the sample graph. This solution contains less than the maximum amount of solute that should dissolve at this temperature and would therefore be classified as a unsaturated solution.
2. A data point corresponding to 70 g KNO₃ per 40 g H₂O (175 g KNO₃/100 g H₂O) and a temperature of 70 °C lies above the best-fit solubility curve in the sample graph. The solution contains more than the maximum amount of solute that should dissolve at this temperature and is supersaturated. See the graph.
   {13966_Answers_Figure_4}
9. Some of the water may have evaporated from the test tubes before their saturation temperatures were measured. What effect would this error have on the solubility of potassium nitrate for a solution? Would the corresponding saturation temperature be too high or too low as a result of this error?

   The concentration in the test tube would be greater than the calculated value because the same amount of solute would be dissolved in a smaller amount of water. The saturation temperature that was measured would be too high for the calculated concentration of the solute.

10. All thermometers have a lag time—it takes a little while to register or report a temperature change. What effect would this error have on the solubility of potassium nitrate for a solution? Would the corresponding saturation temperature be too high or too low as a result of this error?

    The thermometer response time would have no effect on the calculated concentration (solubility) of potassium nitrate. The saturation temperature that would be measured for the calculated solubility would be too high, however, as a result of this error. Remember, the temperature is measured as the solution cools.

References

This experiment has been adapted from Flinn ChemTopic™ Labs, Volume 12, Solubility and Solutions; Cesa, I., Ed., Flinn Scientific: Batavia, IL, 2003.

Introduction

Solubility, defined as the amount of solute that will dissolve in a given amount of solvent, depends on temperature. The solubility of potassium nitrate, for example, increases from 14 g in 100 g of water at 0 °C to about 247 g in 100 g of water at 100 °C—a 1700% increase! While these solubility facts are interesting, they do not allow us to predict the solubility of potassium nitrate at any other temperature. The temperature dependence for the solubility of a substance can only be determined by experiment, by constructing a solubility curve.

Concepts

• Solubility
• Saturated solution
• Saturation temperature
• Solubility curve

Background

A solution that contains the maximum amount of solute that will dissolve at a particular temperature is called a saturated solution. The only practical way to know for sure that a solution is saturated is to observe undissolved solid present in the solution. Undissolved solute plays an active role in the saturated solution. For an ionic compound, ions continually break apart from the undissolved crystal and enter the solution. At the same time, dissolved ions from the solution also recombine to form new crystals. When the solution is saturated, the rate at which the solid dissolves is exactly equal to the rate at which solid recrystalizes from the solution. As a result, the mass of dissolved solute in solution remains constant once the solution is saturated—as long as the temperature does not change. Since the solubility of a substance depends on temperature, the amount of dissolved solute present in a saturated solution also depends on temperature. The solubility of a solute is usually reported as the mass of solute in grams that will dissolve in 100 grams of solvent at a specified temperature. The temperature at which a saturated solution is prepared is called the saturation temperature.

In this experiment, a series of solutions will be prepared, each containing a premeasured amount of potassium nitrate in a known amount of water. The mixtures will be heated to 80–90 °C until all the solid has dissolved. The solutions will then be cooled until the first signs of crystal formation are observed. The temperature at which crystals first appear is the saturation temperature for that concentration of potassium nitrate. The solubility curve for potassium nitrate will be generated by graphing the solubility of potassium nitrate versus the saturation temperature for each solution.

Experiment Overview

The purpose of this experiment is to construct a solubility curve for potassium nitrate in water by measuring saturation temperatures for six different solution concentrations. Working in groups of four, each pair of students will prepare three different solutions and measure their corresponding saturation temperatures. The solubility of potassium nitrate in each solution will be calculated and plotted against the saturation temperature to construct the solubility curve for potassium nitrate in water.

Materials

Prelab Questions

1. Many solutes, including potassium nitrate, have a tendency to remain in solution even after it has been cooled to below the saturation point. This phenomenon is known as supersaturation. Read the entire Procedure carefully. What measure is taken to prevent supersaturation during this experiment?
2. A mixture containing 2.75 g of ammonium chloride (NH₄Cl) in 5.0 g of water was heated to dissolve the solid and then allowed to cool in air. At 61 °C, the first crystals appeared in solution. What is the solubility of ammonium chloride (in g of NH₄Cl per 100 g of water) at 61 °C?
3. The solubility of ammonium chloride in water was measured as described in this experiment and graphed as follows. Use the solubility curve to predict the solubility of ammonium chloride in water at 40 °C.
   {13966_PreLab_Figure_1}
4. One of the students doing the experiment only had time to measure the saturation temperatures for three solutions at 51 °C, 61 °C and 80 °C, respectively. Looking at the graph, do you think the student would have been able to accurately predict the shape of the solubility curve based on these three points? Explain.

Safety Precautions

Potassium nitrate in solid form is a strong oxidant and a fire and explosion risk when heated or in contact with organic materials. It is also a skin irritant. Avoid contact with skin and eyes. Wear chemical splash goggles and chemical-resistant gloves and apron. Use caution when working with a hot water bath and a hot plate. Wear heat-resistant gloves or a heat protector when handling hot glassware. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

1. Form a working group with three other students and divide into two pairs. Each pair of students will prepare three different solution concentrations (A–C) and measure their saturation temperatures. One pair of students will use the Series I masses shown in the reagents table in step 4, the other pair of students will use the Series II masses. Both pairs of students may share the same hot water bath (step 2).
2. Prepare a hot water bath (80–90 °C) for use in step 10: Fill a 250-mL beaker about two-thirds full with hot tap water and place it on a hot plate at a medium-high setting.
3. Obtain three clean and dry test tubes and label them A, B and C. Measure and record the mass of each empty test tube in the data table.
4. Using the following table as a guide, add the recommended amount of potassium nitrate to each test tube A–C. Note that the masses given are ranges—anywhere in this mass range is fine, as long the exact mass used in step 5 is recorded.
   {13966_Procedure_Table_1}
5. Measure and record in the data table the combined mass of each test tube and potassium nitrate.
6. Use a clean, thin-stem pipet to add 20 drops of distilled water to each test tube A–C.
7. Place a 50-mL beaker on the balance pan to support the test tubes. Zero (tare) the balance with the empty beaker in place, then place test tube A in the beaker and measure the mass of the test tube and its contents. Record the total mass of the test tube plus solid plus water in the data table. Note: The mass of water in each test tube should be at least 0.90 g. If not, add one or two more drops of water to the tube and measure the total mass again.
8. Repeat step 7 for test tubes B and C.
9. Place all of the labeled test tubes in the hot water bath.
10. Immerse the thermometer in test tube C. Gently stir the mixture in test tube C using an up-and-down motion of the thermometer until the solid dissolves completely. Note: At this point the solids in the other test tubes should also have dissolved, even without stirring, since they contain less solid.
11. Wearing gloves or using a heat protector, remove the hot water bath from the heat source.
12. Use a test tube clamp to remove test tube C from the hot water bath and place the test tube in the test tube rack and allow the tube and contents to cool slowly in air. Observe the solution closely to watch for the first signs of crystallization.
13. As the solution cools, move the thermometer gently up and down to stir the solution. This will ensure the solution is homogeneous and will maintain a constant, even cooling. Stirring the solution will also encourage crystal formation and prevent supersaturation.
14. Measure and record in the data table the temperature at which the first crystallization begins in the test tube (this is the saturation temperature). Note: The crystals will be colorless in a colorless solution and will not be easy to see. Watch closely—it will look like snow!
15. Remove the thermometer from the test tube and wipe it once with a clean paper towel to remove potassium nitrate crystals. Warm the thermometer briefly in the hot water bath, dry it with a paper towel, and place it in test tube B.
16. Repeat steps 12–15 with test tubes B and A, in that order. Remember to record the saturation temperature for each solution in the data table and to clean the thermometer before transferring it to a new solution.
17. Share the data for the Series I and Series II solutions between both pairs of students in your working group. Complete the data table for all solutions (Series I and Series II).
18. Dispose of the potassium nitrate solutions as directed by your instructor.

Student Worksheet PDF

13966_Student1.pdf