Equilibrium: Solubility

This experiment was developed by Profs LH Berka and NK Kildahl, WPI, based on a commercially available experiment.

1 lab period; work in pairs. Complete the Preparation page before laboratory.

• To discover the use of Le Chatelier's Principle in predicting the effect of temperature on solubility.
• To learn a simple method to estimate the solubility of a salt in water as a function of temperature.
• To learn about solubility product equilibrium, and how the equilibrium constant is related to temperature.

The effect of temperature on the solubility of a solid in water may be predicted using Le Principe du Chatelier if it is known whether the process is endo- or exothermic. If the solution process is endothermic, solubility increases with increasing temperature; if exothermic, solubility decreases. Le Chatelier's Principle permits us to make only qualitative predictions of solubility for a particular solute-solvent system. In this experiment, we will learn how to make our predictions quantitative.

Consider an ionic solid, MX(s), consisting of M⁺ cations and X^- anions, in equilibrium with its ions in aqueous solution:

The equilibrium constant for this process, called the solubility product constant, K_sp, is in (2).

The relationship between K_sp and temperature is in equation (3):

The connection between (3) and Le Chatelier's Principle is DH^o/R, the sign of which depends on whether the dissolution process is endo- or exothermic.

In this experiment, you will determine the water solubility, s (g/mL), of one of the salts, KNO₃, KClO₃, or KBrO₃, as a function of temperature. You will use these data to determine whether dissolution is exo- or endothermic and will determine the quantitative relationship of solubility to temperature expressed in (3).

Focus Questions

1. The procedure used below to measure solution volume is somewhat crude. What are possible sources of error in the procedure?
2. During measurement of solution volume, is it important that the distilled water tube be at the same temperature as the solution tube? Why or why not?
3. How is the solubility of your salt related to its K_sp?
4. For each of your two runs, calculate K_sp for your salt at each temperature, and construct a plot to determine the value of DH^o. Is dissolution exo- or endothermic for your salt?
5. Is the enthalpy of dissolution consistent with predictions based on Le Chatelier's Principle? Explain.

Equipment and Materials

• Labkit
• Iron ring, wire gauze
• 8-inch test tubes
• 600-mL beaker
• Bunsen burner
• KNO₃, KClO₃, KBrO₃

Safety glasses must be worn at all times in the laboratory. You will use mercury thermometers. Mercury is an extremely toxic substance. Report thermometer breakage IMMEDIATELY to your instructor, who will take the necessary steps to clean up the mercury spill. You will work with hot solutions of metal salts; avoid skin contact.

Experimental

Record all data in your notebook. The instructor will assign you one of the three salts to study. Record which salt you are assigned. The following table will guide you in the solution preparation process below.

Set up a 600-mL beaker containing about 500 mL of distilled water on a ring stand, supported over a Bunsen burner on a ring and wire gauze. Heat the water to boiling, then turn off the burner. Relight the burner for brief periods as needed to keep the water fairly hot (> 80 ^oC).

Obtain and clean two IDENTICAL 8-inch test tubes. Rinse them well with distilled water, and dry the insides using a Kimwipe. Use a damp Kimwipe to clean the thermometer.

Carry out the following steps:

1. Weigh a sample of your salt to within 0.1 g of the quantity specified in the table, and transfer it to one of the test tubes. Record the mass of salt. Add to the tube the volume of distilled water specified in the table. Place a thermometer in the test tube, and put the test tube in the hot water bath. NOTE: You will use the thermometer as a stirring rod. STIR GENTLY TO AVOID BREAKING THE THERMOMETER!
2. Stir the salt solution gently until all solid is dissolved. From this point on, do not remove the thermometer from the test tube; you do not want to lose any of the solution.
3. Remove the solution test tube from the water bath, and determine the total volume of the solution by adding water to the second test tube until the level exactly matches the level of solution in the first tube. RAISE THE BOTTOM OF THE THERMOMETER ABOVE THE LEVEL OF LIQUID IN THE SOLUTION TUBE WHILE YOU ARE DOING THIS. Make sure that the solid remains completely dissolved while you are doing this. Return the first tube to the water bath. Transfer the water from the second tube to a graduated cylinder and read its volume to the nearest 0.1 mL.
4. Make sure that all solid in the solution tube is dissolved, then remove it from the bath and allow it to slowly cool in air, with gently stirring, and watch carefully for the onset of crystallization. Record to the nearest 0.1 ^oC the temperature at which solid first appears in the tube.
5. Add an aliquot of water (5 mL for KClO₃, 3 mL for KNO₃ and KBrO₃) to the solution tube, stir, warm if necessary to redissolve the solid, then repeat the previous 2 steps to obtain a second data point.
6. Repeat the previous step at least 3 more times to give a series of points for which the mass of solute is constant, but the volume of solution (hence solute concentration) is changed.

Repeat the above procedure using a second sample of your salt. Adjust the mass of salt so that the saturation temperatures in the second run fall between the saturation temperatures obtained in the first run. You can estimate the required mass of salt as follows. Suppose that in your first run you dissolved 5 g of salt in 5 mL of water (1 g salt/mL water) and observed a saturation temperature of 80 ^oC. You then added 5 mL of water to give a new concentration of 5 g salt in 10 mL water (0.5 g salt/mL water), and observed a saturation temperature of 70 ^oC. To get a saturation temperature of about 75 ^oC will then require a concentration about halfway between those giving saturation T's of 80 and 70 ^oC: 0.75 g salt/mL water. Thus for your second run you should dissolve (0.75 g salt/mL water) * 5 mL water ~ 0.4 g salt in 5 mL of water.

Clean-up. When you have finished all of your work:

• Clean Labkit glassware by recommended procedures, shake off excess water, and return to the Labkit.
• Clean Pasteur pipets, graduated pipets, and vials by recommended procedures, shake off excess water, and place in the drying oven.
• Clean and return all borrowed equipment to the instructor before leaving lab.
• Clean up your work area before leaving lab.

In the event of thermometer breakage, notify the instructor, who will supervise the required disposal procedure. Do not attempt to clean up spilled mercury yourself.

Dispose of salt solutions in the appropriate labelled bottles in the fume hoods.

Preparation
Equilibrium: Solubility