Lab #14: Liquids & Solubility: Cool Solubility Reactions
Prelab Questions:
How can VSEPR models help predict intermolecular bonding? Answer these pre-lab questions BEFORE beginning the
experiment.
1. Predict whether the following molecules are polar or nonpolar. Justify your answer using VSEPR models. Draw them
and fully explain your reasoning!
a) oxygen difluoride, OF2
b) methane, CH4
c) carbon disulfide, CS2
d) fluoromethane, CH3F
e) hydrogen peroxide, H2O2
f) ammonia, NH3
2. As noted by your teacher a couple of minutes ago, the weakest attraction between molecules are collectively called Van
der Waals forces. For each of the above substances, list the kinds of attractive forces between molecules that are expected.
Molecule
LDF
Dipole-dipole
H-Bonds
Oxygen difluoride
methane
Carbon disulfide
Fluoromethane
Hydrogen peroxide
Ammonia
3. What two conditions are necessary for molecules to be polar?
4. If water had a linear molecular shape, would the molecule be polar or nonpolar? Explain your answer.
5. When will hydrogen bonding occur? Give an example of a liquid other than water, in which this type of force is
important.
Part 1: Crystals In and Out
Problem: How does solubility change with temperature?
Materials:
Stirring rod
Thermometer
Large test tube
Metric balance
Potassium dichromate crystals (K2Cr2O7)
Bunsen burner
Ring stand
Hot water bath
Cold water bath
10 ml graduated cylinder
Procedure:
1. Record a hypothesis to answer the problem.
2.
Prepare a hot water bath by using a 400 ml beaker about 2/3 full of water. Place it on the hot plate.
3.
While the water is heating, add between 4 and 5 grams of K 2Cr2O7 to a large test tube. Record the exact mass of the
K2Cr2O7 in the data table below.
4.
Measure out exactly 7.0 ml of water and pour it into the test tube.
5.
Immerse the test tube in the hot water bath.
6.
Insert a thermometer into the test tube and carefully stir with a stirring rod, not the thermometer, until the crystal has
completely dissolved. This will not happen until the mixture becomes relatively hot.
7.
Once the solid has dissolved, remove the test tube from the hot water bath and place it in a test tube rack. Stir gently
as the solution cools.
8.
When the first trace of crystals appear in the test tube, the solution is saturated. Record the temperature at which this
occurs. Save this solution.
9.
Pour exactly 3.0 ml of water into the test tube. The total volume of the solution in the test tube should now be 10
ml.
10. Lower the test tube into the hot water bath and stir until the salt crystals are again completely dissolved in the
solution.
11. Remove the test tube from the hot water bath, stir again, and record the temperature at which the first crystals
appear.
12. Repeat steps 8-10 four more times, adding 3.0 ml of water each time until a total of 22.0 ml of water have been
added. Remember to record the crystallization temperatures in the data table.
13. When you are finished, give the test tube to your teacher for disposal. Clean up your lab area.
Data:
Mass of the K2Cr2O7 =
Total Volume
of Solution
7 ml
10 ml
13 ml
16 ml
19 ml
22 ml
Crystallization
Temperature
Solubility
(g/100 ml H2O)
Questions:
1.
Calculate the solubility of K2Cr2O7 at each of the 6 experimental temperatures. Assume 1 ml of water has a mass of
1 gram. Record the calculated solubilities in the data table above.
2.
On a sheet of graph paper, prepare a plot of the solubility of K 2Cr2O7 versus the temperature. Be sure to place the
dependant and independent variables on the appropriate axes. Include a title on your graph.
3.
What does your graph indicate about the relationship between solubility and temperature?
4.
Look up the solubility of K2Cr2O7 at various temperatures in the Handbook of Chemistry and Physics. Enter these
accepted values on your solubility graph and draw a curve connecting the values using a different color pen. Be sure
to includes a key on your graph indicating which line represents your experimental data and which is the plot of the
actual curves.
5.
How does the accepted solubility curve compare to your experimental solubility curve?
6.
Name several sources of error that might explain any differences in the curves.
Part 2: Yellow Trouble
Problem: How do the concentrations of the reactants affect the amount of precipitate formed in a reaction?
Materials:
Potassium iodide crystals (KI)
Lead nitrate crystals (Pb(NO2)2)
25 ml graduated cylinders
10 ml graduated cylinders
Stirring rod
Metric balance
4 test tubes
Test tube rack
Introduction:
It’s time to put your new solution calculation skills to the test. When a chemical solution is prepared by your teacher, the
proper calculations must first be completed. In this exploration experiment, you will be taking on the job of the teacher by
determining the amounts of solid that must be weighed out to prepare different solutions.
Hypothesis:
Before actually performing the experiments, look closely at the ingredients for the four reactions. A solid, PbI2, will form
when the 2 chemicals combine. Which reaction do you expect to produce the largest amount of solid? Explain your
prediction.
Procedure:
1.
Do all the calculations necessary to determine the number of grams of each solute needed to prepare the solutions
with their given concentrations. Show all of your calculations, including units.
Solution 1: 20 ml of 0.2 M KI
Solution 2: 20 ml of 0.2 M Pb(NO2)2
Solution 3: 20 ml of 0.4 M KI
Solution 4: 20 ml of 0.4 M Pb(NO2)2
2.
Once your teacher has checked your calculations, you may proceed to prepare the four solutions. Be extremely
careful in weighing out your solids because you will be measuring fairly small amounts. You will be graded on the
accuracy of your results.
3.
Pour each of the solutions into four separate large test tubes and place them in a test tube rack. Label the solutions
according to their concentrations and content.
4.
Now, use these solutions to carry out the four chemical reactions outlined in the table below. Combine the proper
portions of each chemical in a small graduated cylinder. Make sure to wash the cylinder between reactions.
5.
Do not throw your precipitate down the drain or in the garbage. Your teacher will show you where to properly
dispose of your chemicals.
6.
Clean up your lab area.
Data:
Reaction
Letter
A
B
C
D
Amount of KI
Amount of Pb(NO2)2
10 ml of 0.4 M KI
10 ml of 0.4 M KI
10 ml of 0.2 M KI
10 ml of 0.2 M KI
10 ml of 0.4 M Pb(NO2)2
10 ml of 0.2 M Pb(NO2)2
10 ml of 0.4 M Pb(NO2)2
10 ml of 0.2 M Pb(NO2)2
Height of the Precipitate (ml)
Questions:
1.
Write a balanced equation for the reaction that occurred above. Keep in mind that the solid formed is lead iodide.
2.
Using the data above, determine the limiting reagents in each equation (if applicable), and the amounts of each
product produced.
3.
Which of the reactions formed the largest amount of precipitate?
4.
Was the prediction you made correct?
5.
Explain the similarities in the heights of the precipitates formed in reactions A and B.
Part 3: A Salty Dilemma
Problem: How can you determine the solubility of sodium chloride?
Materials:
Saturated salt solution
Bunsen burner
Watch glass
Thermometer
Ring stand
Tongs
Evaporating dish
Procedure:
1. Record a hypothesis to answer the problem.
2.
Measure out approximately 20 ml of saturated salt solution.
3.
Take the temperature of the solution and record it in the data table below.
4.
Weigh a dry, clean evaporating dish and a large watch glass and record it in the data table below.
5.
Pour the measured salt solution into the evaporating dish. Cover it with the watch glass.
6.
Weigh the dish, glass and solution. Record the weight in the data table below.
7.
Place the evaporating dish on the ring stand.
8.
Heat the solution gently, being careful not to let the solution boil too vigorously. If you allow your solution to boil
over, you will need to start again. Low heat for about 10 minutes should be enough to evaporate all the liquid.
9.
When evaporation is complete, allow the dish and glass to cool completely.
10. Weigh the dish, glass, and solid. Record the weight in the data table below.
11. Calculate the solubility of the salt from your data.
Data:
Temperature of the solution
Mass of the evaporating dish and watch glass
Volume of the liquid
Mass of the dish, glass, and solution
Mass of the dish, glass, and solid after evaporation
Mass of the solid (salt)
Solubility of NaCl (g/100 ml H2O)
Questions:
1.
What is the accepted value of solubility of NaCl in water? (Use the Handbook of Chemistry and Physics.)
2.
How does your experimental solubility compare to the accepted value (is it too high or too low or about the same)?
3.
Explain and differences you observed between your experimental and accepted solubility values. (What may have
gone wrong?)
Part 4: Soda Solution
Problem: What is the concentration of carbon dioxide gas in soda?
Materials:
Soda
Metric balance
Graduated cylinders
Beakers
Hot Plate
Introduction:
Everyone has had soda. Have you ever given much thought to what is in it? Soda has a fairly simple composition: water,
sugar, flavoring, and carbon dioxide are the main ingredients.
Procedure:
Record a hypothesis to answer the problem. Your challenge is to design an experiment that will allow you to determine the
concentration of carbon dioxide in a sample of soda. Write your procedure below. Have your teacher check your idea before
you begin collecting your data. Remember, your goal is to be able to calculate the concentration of CO 2 in your soda sample.
Make certain you have collected the right kind of data to do this. Create a table for your data.
Questions:
1.
Calculate the molarity of CO2 in soda. Show all of your calculations.
2.
Using your data, determine the mass of CO2 in a 2 L bottle of soda. Show your work.
3.
Do you think your CO2 calculation is truly representative of the actual mass of CO 2 in the soda? Why or why not.
4.
Check with another person in the lab who used a different kind of soda. Did you get the same results? Why or why
not?
Part 5: Milky Solutions
Overview:
Oil causes the food coloring to circulate in the milk making swirls of color
Equipment:
Milk (in a cup or single serving carton with the lid cut off)
Food colors
vegetable oil
a table to place the milk on
Procedure:
1. Put the milk in a cup, about 1/4 full. 2% or higher fat content works the best.
2. Place one drop of each color of food coloring in each corner of the carton or near the sides of the glass.
3. Place a drop of vegetable oil in the middle of the cup. Be careful not to move the glass or to shake the table. After a
short time the colors swirl and mix, it looks really cool.
4. Try it with different fat percentages and see how it affects the patterns or the rate of swirling.
5. Record your observations below.
Data and Analysis:
Type of Milk
Swirl Pattern
Time to Swirl
Skim Milk
1% Milk
2% Milk
Whole Milk
Questions:
1.
Which milk took the longest time to swirl? The shortest?
2.
What are the differences between the 4 types of milk?
3.
Why does the food coloring swirl like this in milk? Use the terms miscible and immiscible in your explanation.
4.
Do some research…is milk good for you or not? Support your answer.
Part 6: Polarity & Solubility
Molecules with polar bonds can be polar or non-polar. If they are completely symmetrical, the symmetry of the
molecule can cancel out the polarity of the bonds. If a molecule is polar, it may have greater attraction for other
polar molecules than non-polar molecules will. Non-polar molecules have an even distribution of charge, and will
not attract one another, and thus will mix quite well. In general, the phrase “ Like dissolves like” describes the
solubility of solutes in solvents. That is, polar solvents such as water, will easily dissolve other polar molecules;
polar solvents will also dissolve ionic compounds because ionic compounds represent the extreme case of a polar
molecule. Non-polar solvents will generally dissolve non-polar solutes.
In this activity, you will determine the solubility of a variety of solutes in two different solvents, water and hexane.
Procedure:
1. Obtain 7 test tubes of the same size, and place 5 mL of water in each one. (After you have measured the first one,
you may estimate the 5 mL level in the other 6 tubes.)
2. Test the water solubility of each of the seven solutes by adding a different solute to each test tube. For liquid solutes,
add about 20 drops. Transfer a pea-sized sample of each solid solute using a spatula or a wood splint. Gently mix
the test tube and contents by flicking the bottom with your finger.
3. Record your observations in your lab book: S = soluble, SS = slightly soluble, IN= insoluble.
4. Discard the samples into the waste container labeled “Water Waste.”
5. Clean and dry the test tubes and repeat using hexane as the solvent.
6. Dispose of the hexane and solute wastes in the container labeled “Hexane Waste.”
7.
Solute
Water ( H2O)
Hexane (C6H14)
Urea ( CO(NH2)2 )
Iodine ( I2)
Ammonium chloride
Naphthalene ( C10H18)
Copper(II) sulfate
Ethanol ( C2H5OH)
Sodium chloride
Questions:
1. Which solutes were more soluble in water than in hexane? Why?
2.
Which solutes were more soluble in hexane that in water? Why?
3.
What can be said about the polarity of each of the solutes?
Part 7: Capillary Action
Capillary action depends partially on the polarity of the molecules of the liquid and of the material that forms the
tube or space into which the liquid rises. A polar attraction between the material of the tube and the molecules of a
polar liquid causes the liquid to rise into the tube. As one molecule moves up, it attracts neighboring molecules
which, in turn attract their neighbors. Once the upward attractive force is equal to the downward force of gravity,
the liquid stops rising. An extreme example of this phenomenon occurs in giant trees like redwoods. It is estimated
that these large trees move about 2000 liters of water each day from their roots to the uppermost leaves by capillary
action.
Procedure:
1. Obtain a well-plate, eight capillary tubes, and a container of water.
2. Place a few drops of water in one of the wells in the plate.
3. With one end of the tube, touch the water. The top of the tube must be open. Leave the tip of the tube in contact with
the liquid for at least 15 seconds.
4. Once the liquid stops rising, remove the tube from the well and measure the height of the column of liquid. Record
the value.
5. Drain the capillary tube by touching its tip to a paper towel.
6. Repeat the procedure to allow liquid to rise into the tube 2 more times. Record the heights in the table.
7. Repeat the procedure using different liquids in the wells, and a new capillary tube for each liquid. (Note: the
glycerol will take more than 15 seconds to reach its final height.)
8. Calculate the average height of the liquid columns for each substance tested.
Liquid
Water
Ethanol
Ethylene glycol
Glycerol
Mineral Oil
Propanol
Salt Water
#1 ( Height)
#2
#3
Average
Based on the heights of the liquid columns, draw a conclusion about the relative polarity of the solvents tested.
Part 8: Introduction to Intermolecular Forces in Solution
Hydrogen bonding is due to the polarity of the oxygen-hydrogen, nitrogen-hydrogen, or fluorine-nitrogen bonds in
molecules. Hydrogen bonding is an intra-molecular bond. That means it occurs between molecules, not within a
molecule. Hydrogen bonding in water is responsible for some of water’s interesting properties. One of these
properties is surface tension. In this activity, you will investigate hydrogen bonding and surface tension by seeing
how many drops of a liquid you can place on a penny before it runs over.
Trial number
Number of drops
Trial #1
Trial #2
Trial #3
average
You will be given three solutions: sodium carbonate, sodium chloride and a 1% liquid detergent sample. Your task
is to design and carry out an experiment to determine their effect on the surface tension of water.
Design a procedure which includes a control, steps that will ensure that the results are reproducible, and predictions
of the effect of each solute.
Create a table that will organize your data clearly.
Questions:
1. Compare the average number of drops placed on your penny with the results obtained by some of the other lab groups.
What might account for differences?
2.
What were your predictions for the effect of each solute on the surface tension of the water?
3.
What was your control?
4.
What steps did you include in your procedure to increase the reliability or reproducibility of your results?
5.
Will the addition of salts always change the surface tension of water? Why or why not?
6.
What effect did the detergent have on the surface tension?
7.
What are some practical applications of what you tested and learned?