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# 19.4: Slightly Soluble Salts

Difficulty Level: At Grade Created by: CK-12

## Student Behavioral Objectives

The student will:

• define solubility product constants.
• write solubility product constant expressions.
• calculate solubility product constants.

## Timing, Standards, Activities

Timing and California Standards
Lesson Number of 60 min periods CA Standards
Slightly Soluble Salts 1.5 None

### Activities for Lesson 4

Laboratory Activities

1. None

Demonstrations

1. None

Worksheets

1. Solubility and Solubility Product Constant Worksheet

1. None

## Answers for Slightly Soluble Salts (L4) Review Questions

• Sample answers to these questions are available upon request. Please send an email to teachers-requests@ck12.org to request sample answers.

## Multimedia Resources for Chapter 19

The two videos below provide an introduction to chemical equilibrium.

This website provides a video on Le Chatelier’s Principle.

This webiste provides a video on Le Chatelier’s Principle and the Haber Process.

A Khan Academy electronic lecture on Le Châtelier's Principle is available at

## Teacher’s Resource Pages for Solution Equilibria

Investigation and Experimentation Objectives

In this activity, the student will, through a series of laboratory activities, develop evidence for the existence of the equilibrium state, and answer a series of questions to develop a conclusion about the nature of the equilibrium state.

Lab Notes:

This lab involves the $Fe^{3+}/SCN^-$ equilibria, where the $Fe[SCN]^{2+}$ complex is formed. The value of K at room temperature is about 138. The equilibrium original equilibrium position produced by the lab directions will be a light reddish orange and this position will be shifted by various activities during the lab.

Solution Preparation

To prepare $250.~ mL$ of $0.10~M~FeCl_3$ solution from solid anhydrous iron (III) chloride, dissolve $4.1~ g ~FeCl_3$ in a $250.~ml$ volumetric flask and fill to the line. If you are using iron (III) chloride hexahydrate, dissolve 6.8~g~ FeCl_3 \cdot H_2O[/itex] in a $250.~ml$ volumetric flask and fill to the line.

To prepare $250.~ mL$ of $0.10~ M~KSCN$ solution from solid $KSCN$, dissolve $2.4~g$ of the solid in a $250.~ml$ volumetric flask and fill to the line.

The following questions refer to a hypothetical reversible chemical reaction in which reactant Y is a bright yellow color, reactant C is colorless, and the product B is a bright blue color.

$C ~+~Y~ \leftrightharpoons ~3B$

1. At the original equilibrium position, the solution is green. Describe the color change you would expect if a large quantity of C were added to the reaction solution.

Since the addition of a large quantity of C will drive the equilibrium to the right, yellow color (Y) will decrease significantly and blue color (B) will increase . . . so the final solution will change from green to blue.

2. At the original equilibrium position, the solution is green. Describe the color change you would expect if a large quantity of C were removed from the reaction solution.

Since the removal of C from the solution would shift the equilibrium to the left, the blue color (B) would decrease and yellow color (Y) would increase . . . so the final solution would change from green to yellow.

1. When $FeCl_3$ was added to the equilibrium solution in test tube 2, which way did the equilibrium shift? How do you know?

Forward . . . darker red color indicated more product was formed.

2. When $KSCN$ was added to the equilibrium solution in test tube 3, which way did the equilibrium shift? How do you know?

Forward . . . darker red color indicated more product was formed.

3. When crystals of potassium phosphate were added to the equilibrium solution in test tube 4, which way did the equilibrium shift? How do you know?

Backward . . . lessening of red color indicated product was being used up.

4. If more product formed in test tube 2 when $FeCl_3$ was added, what other substance had to be present in the equilibrium solution before the $FeCl_3$ was added? How do you know?

$KSCN$ . . . in order to form more product, both reactants had to be available.

5. If more product formed in test tube 3 when $KSCN$ was added, what other substance had to be present in the equilibrium solution before the $KSCN$ was added? How do you know?

$FeCl_3$. . . in order to form more product, both reactants had to be available.

6. Therefore, what substances were present in the equilibrium solution in test tube 1?

$Fe^{3+}~~~\text{and}~~~SCN^- ~~~\text{and}~~~Fe[SCN]^{2+}$

7. Given the information that phosphate ion removes iron (III) ion from solution, describe what happened when phosphate ion was added to test tube 4.

The added phosphate ions removed $Fe^{3+}$ causing the equilibrium to shift backward, thus using up $Fe[SCN]^{2+}$, resulting in a decrease in the red color.

## Solution Equilibria Laboratory

Background Information

A state of equilibrium is affected by concentration of reactants, temperature, and pressure (for reactions containing gaseous substances). If a system at equilibrium is subjected to a change in one or more of these factors, a stress is placed on the equilibrium. When a stress is places on a system at equilibrium, the equilibrium position will shift in the direction that tends to relieve the stress.

Materials and Apparatus (per lab group)

Beaker, 100 mL Graduated cylinder, 10 mL Test tubes (4), small Test tube rack Dropper pipet 0.10 M $FeCl_3$ solution, 10 mL 0.10 M $KSCN$ solution, 10 mL Solid potassium or sodium phosphate, a few crystals Distilled water

Safety Issues

Safety glasses and apron should be worn at all times while working in the chemistry laboratory.

Procedure

$Fe^{3+}~+~SCN^-~ \leftrightharpoons ~ Fe[SCN]^{2+}$

The iron (III) ion solution is light brown, the thiocyanate ion solution is nearly colorless and the product complex is dark red.

1. Measure 5 mL of $0.10~M~FeCl_3$ solution and place it in a 100 mL beaker. Add 5 mL of $0.10~M~KCSN$ to the same beaker. Dilute the contents of the beaker with distilled water until the solution is a light reddish-orange color. Divide the solution equally among the four numbered test tubes. Set tube 1 at one end of the rack to be used for color comparison.

2. Using a dropper pipet, add $0.10~M~FeCl_3$ solution drop by drop to the solution in test tube 2 with stirring until a significant color change occurs. Record your observations and rinse the pipet with distilled water.

3. Add $0.10~M~KCSN$ solution drop by drop to the solution in test tube 3 with stirring until a significant color change occurs. Record your observations and rinse the pipet with distilled water.

4. Drop a couple of crystals of potassium phosphate into the solution in test tube 4 with stirring. Record your observations.

Data:

Data Table
Test Tube Number Color
1
2
3
4

Post-Lab Questions

1. When $FeCl_3$ was added to the equilibrium solution in test tube 2, which way did the equilibrium shift? How do you know?

2. When $KSCN$ was added to the equilibrium solution in test tube 3, which way did the equilibrium shift? How do you know?

3. When crystals of potassium phosphate were added to the equilibrium solution in test tube 4, which way did the equilibrium shift? How do you know?

4. If more product formed in test tube 2 when $FeCl_3$ was added, what other substance had to be present in the equilibrium solution before the $FeCl_3$ was added? How do you know?

5. If more product formed in test tube 3 when $KSCN$ was added, what other substance had to be present in the equilibrium solution before the $KSCN$ was added? How do you know?

6. Therefore, what substances were present in the equilibrium solution in test tube 1?

7. Given the information that phosphate ion removes iron (III) ion from solution, describe what happened when phosphate ion was added to test tube 4.

## A Light Activated Reversible Chemical Reaction

What the student sees.

The room is somewhat darkened. The students see a large beaker containing a purple solution sitting on an overhead projector. Half of the glass plate of the overhead project is covered with several folded layers of aluminum foil. The beaker is sitting such that half of it is on the aluminum foil and half of it is on the glass plate. When the overhead projector is turned on, the half of the solution in the beaker that is above the light turns clear . . . the other half remains purple.

The solution appears to be divided by an invisible line running vertically through the beaker. When the overhead projector is turned off, the colorless side of the solution will slowly return to purple. The reaction can be repeated and is reversible for a couple of days.

Procedure:

1. Mix together in a 1-liter beaker

• $10.\ mL$ of freshly prepared $0.001\ M$ thionin solution. $0.001\ M$ thionin solution can be prepared by dissolving $0.023 \ g$ of Thionin in $100.\ mL$ of distilled water. (Thionin solution has a shelf life of only a few days – so prepare it only when you are ready to use it.)
• $100.\ mL$ of $1.0\ M$ $H_2SO_4$
• sufficient distilled water to bring the total volume to $600\ mL$.
• mix thoroughly

2. Turn off the room lights and add $2.0 \ grams$ of iron(II) sulfate. Stir to dissolve.

Disposal

The solution can be rinsed down the drain followed by excess water.

Discussion

The two forms of thionin differ in oxidation state and the redox reaction converting one form to the other also involves the conversion $Fe2+$ ions to $Fe3+$ ions. The activation energy for the reaction is provided by light.

Source of materials

Most chemical supply companies carry all the materials necessary for this demonstration. A kit specifically designed for this demonstration is offered by Flinn Scientific.

## Equilibrium Between Nitrogen Dioxide and Dinitrogen Tetroxide Demonstration

Apparatus and Materials:

• Three sealed borosilicate glass tubes filled with nitrogen dioxide gas - demo tubes are available from most chemical supply companies (store in bubble wrap)
• Three tall form 1-liter beakers
• Hot plate
• Tap Water
• Ice

Procedure:

1. Fill one of the beakers with tap water and place it on a hot plate, heat until boiling.
2. Fill another beaker with a mixture of ice and tap water.
3. Displace the sealed tubes containing nitrogen dioxide (should be same color due to same temperature). A white background makes them much more visible.
4. Place one tube in each of the hot and cold beakers and leave the third at room temperature.
5. As the equilibrium in the tubes adjust to the new temperatures, there will be significant differences in the colors of the tubes. The hot tube becomes dark brown, the cold tube may become nearly colorless.
6. The tubes can be removed from the beakers and switched to the other beaker and the colors will change again.

Safety Issues:

The borosilicate tubes can easily withstand the temperature extremes of these two water baths, care should be taken, however, to not damage the tubes and allow the nitrogen dioxide to escape. Nitrogen dioxide is an extremely toxic gas. It is irritating to the respiratory tract.

Discussion:

The equilibrium illustrated in this demonstration is between nitrogen dioxide, $NO_2$ and dinitrogen tetroxide, $N_2O_4$. The chemical equation is shown below.

$2 \ NO_{2(g)} \leftrightharpoons \ N_2O_{4(g)}$

Nitrogen dioxide is a dark reddish brown gas and dinitrogen tetroxide is a colorless gas. When the equilibrium is shifted to the left, as written above, the amount of nitrogen dioxide increases, the amount of dinitrogen tetroxide decreases, and the color of the tube darkens. When the equilibrium is shifted to the right, as written, the amount of nitrogen dioxide decreases, the amount of dinitrogen tetroxide increases, and the color of the tube lightens. As written, the reaction releases $58 \ kJ$ of energy. Since this is an exothermic reaction, increasing the temperature will drive it to the left and decreasing the temperature will drive it to the right.

Aug 18, 2012

Aug 13, 2014