CH105
Pre-Lab 4:
Solutions
Directions: Read the Goals, Background, Safety, and Procedure sections for this
experiment, then answer the following questions in the space provided. For short
answer questions, write complete sentences and provide a reason for the answer. For
calculation questions, show all work and report answers in a box with the appropriate
significant figures and units. Pencil is acceptable for this assignment.
1. According to Figure 1 in the lab, what is the solubility of
NaCl in water at 100˚C in grams of solute per 100 g H2O?___________________
2. According to Figure 1 in the lab, what is the solubility of
NH4Cl in water at 100˚C in grams of solute per 100 g H2O?_________________
3. According to Figure 1 in the lab, at room temperature
(~25˚C), what is more soluble in water NaCl or NH4Cl?______________________
4. What is the molarity of a 275 mL solution that contains 32.8 g of sodium
chloride?
5. What is the mass percent of sugar in 368 g of solution that contains 24.5 g
of sugar?
6. What should dissolve faster: a chicken
bouillon cube (with a volume of 1 teaspoon)
or a teaspoon of chicken bouillon granules?________________________________
(Chicken bouillon is evaporated seasoned meat extract that is available at your
grocery store.) Explain.
7. What instructions would you write on the chicken bouillon package to help
consumers easily dissolve the chicken bouillon?
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PCC/Rock Creek
Experiment 4:
Solutions
Goals: • Test the relationships between solubility and temperature as
well as rate of dissolving and particle size for solid solutes in
water.
• Prepare unsaturated, saturated, and supersaturated solutions.
• Examine the properties of unsaturated, saturated, and
supersaturated solutions.
• Calculate the concentration of solutions in terms of molarity
and mass percent.
• Qualitatively describe the solubility of solutes in various
solvents.
• Test the conductivity of several solutions.
Purpose: Investigate the formation and properties of solutions, factors
affecting solubility, and calculation of concentration.
Background:
Matter is often divided into two major categories: pure substances and
mixtures. Pure substances are elements and compounds. Mixtures, on
the other hand, contain 2 or more pure substances. These materials can
be evenly distributed throughout the entire mixture (homogeneous
mixture) or unevenly distributed throughout the entire mixture
(heterogeneous mixture). Solutions are an example of homogenous
mixtures.
A solution is composed of a solvent and a solute. The solvent is largest
component of the mixture. The solute is the smaller component of the
mixture. In an aqueous salt solution, water is the solvent and a salt is
the solute. (Hint: Solvent has 7 letters and solute has 6 letters. The
longer word corresponds to the larger component!)
Solute-Solvent Systems
A solution can be made from a variety of solute-solvent systems. Matter
comes in three phases (solid, liquid, gas) and the components of a
solution can be in any of the phases of matter. For this laboratory
experiment, you will focus on the solid solute/liquid solvent and liquid
solute/liquid solvent systems.
A general rule for solubility is “like dissolves like.” This expression
implies that polar solvents will usually dissolve polar solutes, while
nonpolar solvents will usually dissolve nonpolar solutes. Polar solvents
do not usually dissolve nonpolar solutes and nonpolar solvents do not
usually dissolve polar solutes.
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PCC/Rock Creek
Solutions • 3
The intermolecular forces at work between a solvent and a solute in a
solution have to be great enough to overcome the attraction the solute
molecules have for each other as well as the attraction the solvent
molecules have for each other. While “like dissolves like” provides
general guidelines for solubility, the solubility rules for ionic compounds
as well as the solubility table in your book provide more specific
information about ionic compound solubility.
Factors Affecting Solubility
The solubility of a solute in a solvent depends not only on the chemical
nature of the solute-solvent system, but it also depends on the physical
natures of the solute and solvent. Factors that affect solubility include:
temperature, particle size, agitation, and concentration of a solution.
As a general rule, solid solutes tend to become more soluble in liquid
solvents as the temperature of the solvent increases. (See Figure 1
below.) The opposite trend exists for gases dissolved in a liquid. A
carbonated beverage at room temperature is never as fizzy as a
carbonated beverage right out of the refrigerator because the gas escapes
at warmer temperatures.
Figure 1. Solubility Curve.
PCC/Rock Creek
4 • Experiment 4
Smaller particles dissolve better than larger particles. Small particles
have more surface area for the solvent molecules to attack than larger
particles do. Consequently, the solvent molecules are able to dissolve
small particles more easily than large particles of the same material.
Agitation also influences the solubility of a solute. A solute that is
poured into a solvent without stirring is less likely to dissolve than the
same solute that us pour into a solvent and stirred. By stirring the
solute particles, the solvent makes better contact with the solute than
without stirring. Consequently, the solution forms more rapidly when
the solution is agitated.
Solute dissolves more readily in a dilute solution than a concentrated
solution. In a dilute solution, a larger number of solvent molecules are
attracted to the solute particles because of the large differences in ratios
between the number of solute and the number of solvent molecules.
Description of Solubility
The solubility of a solute can be described in a number of ways. The two
broad categories for these descriptions are: qualitative and quantitative.
A qualitative description of solubility is a more general (less precise)
description than a quantitative description. A qualitative description of a
solute involves its classification as soluble (dissolves readily), insoluble
(does not dissolve), or slightly soluble (a small amount dissolves) in a
specific solvent. A quantitative description of solubility is more specific
than the qualitative description and frequently involves the number of
grams of solute that are soluble in a specific number of grams (or
milliliters) of solvent.
Degree of Saturation: Qualitative Description of Solutions
A qualitative description of a solution involves the classification of the
solution as saturated, unsaturated, or supersaturated. A saturated
solution is a solution in which no more solute will dissolve in the solvent.
An unsaturated solution, on the other hand, is a solution in which more
solute will dissolve in the solvent. A supersaturated solution is a
solution in which more solute than should dissolve in the solvent is
dissolved at a given temperature.
Supersaturated solutions are made by slowly cooling a saturated
solution at a high temperature. Because the solubility of solids generally
decreases as temperature decreases, a solution that is saturated at a
high temperature is supersaturated at a lower temperature. These
solutions are highly unstable. Shaking, adding a seed crystal, or
scratching the glass inside the container of a supersaturated solution
may cause crystallization to occur.
Solutions • 5
Concentration: Quantitative Description of Solutions
The concentration of a solution can be expressed in several different
ways. Among these expressions are: molarity, molality, percent mass,
percent volume, parts per million, and parts per billion. In this
experiment, you will calculate molarity and percent mass for two
solutions.
Molarity is the ratio of moles of solute to the volume of solution in liters.
The equation is shown below:
Molarity 
moles of solute
liters of solution
Mass percent is the ratio of the mass of the solute to the mass of the
solution multiplied by 100%. The equation is shown below:
Mass percent 
mass of solute
mass of solution
100%
Conductivity
You will also measure the conductivity (ability to pass an electrical
current) of several solutions. Solutions that conduct electricity contain
electrolytes. Solutions that do not conduct electricity contain
nonelectrolytes. Electrolytes are usually ionic compounds that
disassociate in solution. The “free ions” are able to carry the electrical
current through the solution.
Procedure:
Solubility and Temperature
1. Label one test tube NaCl and another NH4Cl.
2. Place 1.0 g NaCl into the NaCl tube and 1.0 g NH4Cl into the NH4Cl
tube.
3. Add 5 mL of room temperature water. Shake gently until both salts
are dissolved.
4. Add a tiny crystal of the appropriate compound and observe any
changes. Record these on your data sheet.
5. Add 1.4 g more NaCl into the NaCl tube and 1.4 g more NH4Cl into
the NH4Cl tube. Shake gently for about 5 minutes and note the
results on your data sheet.
PCC/Rock Creek
6 • Experiment 4
6. Place both tubes into a beaker of gently boiling water and periodically
shake the tubes. After 5 minutes of heating, note the results on your
data sheet.
7. Place the warm tubes in a test tube rack and allow them to cool
(about 10 minutes). Note any changes. If no change has occurred in
the NH4Cl tube, gently scratch the inside of the tube with a stirring
rod. Note the results on your data sheet.
Unsaturated, saturated, supersaturated
1. Place 5 mL of distilled water into a 50 or 100 mL beaker. Record the
temperature of the water on your data sheet.
2. Mass ten (10) individual 0.25 g samples of KCl.
3. While stirring with a glass rod, add KCl to the water in 0.25 g
portions; keep adding until no more KCl dissolves. The solutions
should be saturated. Record the mass of the KCl added on your data
sheet.
4. Transfer the saturated KCl solution to a graduated cylinder. Record
the volume of the solution on your data sheet.
5. Calculate the molarity of a saturated KCl solution at room
temperature.
6. Transfer the saturated KCl solution back into the beaker.
7. Add the remaining 0.25 g samples to the KCl and stir. Notice the
additional KCl does not dissolve.
8. Heat the solution on a hot plate until all of the solid dissolves. With a
hot glove or tongs, remove the beaker from the hot plate and set it on
the bench top, out of the way.
9. Place an applicator stick, or suspend a string, into the solution and
allow the solution to cool. Continue with the next part of this
experiment and return to this part after the solution cools to room
temperature.
10. Observe what happened to the solution when it cooled to room
temperature. Offer an explanation for what has taken place. (If no
crystals have formed, drop into the solution a single crystal of KCl or
stir the solution with a stirring rod.)
Solutions • 7
Percent sodium chloride in a brine solution
1. Select an evaporating dish and a beaker from your drawer that the
evaporating dish will rest securely on. (This is usually a 250-mL
beaker!)
2. Clean your evaporating dish and heat it for 2 minutes on the hot plate
with gentle heat.
3. Allow the evaporating dish to cool, then obtain the mass of the
evaporating dish and record the value on your data sheet.
4. Fill the beaker approximately ¾ full with water. Heat it on the hot
plate.
5. Place 6-10 mL of brine solution in the evaporating dish and obtain the
mass of the evaporating dish and solution. Record this value on your
data sheet.
6. Carefully place the evaporating dish on the partially-filled beaker and
boil the water until the material in the evaporating dish looks dry
(about 30 minutes).
7. Remove the evaporating dish from the beaker and place it directly on
the hot plate. Remember, hot glass looks like cold glass at these
temperatures. You probably will want to use a hot pad or glove.
8. Heat the evaporating dish on the hot plate for 1-2 minutes. Watch for
“popping” sodium chloride.
9. Remove the evaporating dish from the hot plate and allow it to cool.
10. Mass the evaporating dish with the solid and record this value on
your data sheet.
11. Reheat the sample, cool, and remass the sample until a constant
mass is obtained.
12. Calculate the percentage of NaCl dissolved in a saturated brine
solution at room temperature.
PCC/Rock Creek
8 • Experiment 4
Polar or Nonpolar
(This experiment can be divided into 2 parts. One team can perform the experiment with distilled water and
ethanol, while the other team performs the experiment with acetone and petroleum ether. Each team must
make their own observations.)
1. Obtain 16 clean, dry small test tubes.
2. Make 4 sets of 4 test tubes each containing approximately 0.1 g of
sodium chloride (NaCl), sucrose (C12H22O11), naphthalene (C10H8), and
iodine (I2). (Do not mass these substances, your instructor will have examples of what 0.1 g looks
like for each substance.)
3. Add 3 mL of distilled water to each test tube in the first set and tap
the test tube with your finger to mix the contents. Record whether
the solid dissolved completely (soluble), partially (slightly soluble), or
not at all (insoluble).
4. Check the conductivity for each of the solutions and record your
findings on your data sheet.
5. Repeat steps 3 and 4 with ethanol, acetone, and petroleum ether
instead of water. Record your observations on your data sheet.
6. The water solutions should be placed into the aqueous waste
container and the solutions from step 5 should be placed into the
organic waste container.
Rate of dissolving vs. particle size
1. Place 0.5 g small crystalline NaCl in one large test tube and 0.5 g
large crystalline NaCl in another large test tube.
2. Add 20 mL of water to each tube and shake gently.
3. Record the time required to dissolve each amount of NaCl.
Name:______________________________
Date:_____________
CH105/Sp06
Data Sheet 4:
Solutions
Directions: Record the data as it is collected onto this sheet in BLUE or BLACK ink. Do
not use white out. Correct mistakes by making a single line through the error and
writing the new information above or beside the mistake.
Solubility and Temperature
Time
NaCl Observations
After small
crystal is
added to the
test tubes.
After 1.4 g
solid added
and shaken
gently for 5
minutes.
After placed
in gently
boiling water
for 5 min.
and shaking
periodically.
After allowed
to cool for 10
minutes.
NH4Cl Observations
Unsaturated, saturated, supersaturated
Temperature of water (˚C)
________________
Mass of KCl added to water (g)
________________
Moles of KCl added to water
________________
Volume of the solution (mL)
________________
Molarity of the KCl solution (M)
________________
Observations after solution is allowed to cool:
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PCC/Rcok Creek
Experiment 4
Data Sheet
Percent sodium chloride in a brine solution
Mass of the evaporating dish (g)
_________
Mass of evaporating dish + brine solution (g)
_________
Mass of brine solution (g)
_________
Mass of evaporating dish + dried salt (g)
_________
Mass of dried salt (g)
_________
_________
_________
Mass percent of salt in brine solution at room temperature _________
Polar or Nonpolar
Solubility
Solvent/Solute
water
ethanol
acetone
petroleum ether
Conductivity
Solvent/Solute
water
ethanol
acetone
petroleum ether
NaCl
C12H22O11
C10H8
I2
NaCl
C12H22O11
C10H8
I2
Rate of dissolving vs. particle size
Material
Time/Observations
small NaCl
large NaCl
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PCC/Rock Creek
Name:______________________________
Date:_____________
CH105/Sp06
Post Lab 4:
Solutions
Directions: Answer the following questions in the space provided. For short answer
questions, write complete sentences and provide a reason for the answer. For
calculation questions, show all work and report answers in a box with the appropriate
significant figures and units. Pencil is acceptable for this assignment.
1. Candle wax is nonpolar. Would you
expect candle wax to dissolve in water?_____________________________
Explain why or why not.
2. Candle wax is nonpolar. Would you
expect candle wax to dissolve in petroleum ether?____________________
Explain why or why not.
3. Is sugar a good electrolyte?__________________________________________
Explain why or why not.
4. Does crushed ice melt faster or slower than
a large ice cube with the same volume in water?______________________
Explain your reasoning.
5. When the ice melts in question 4, does a solution form?______________
Explain why or why not.
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PCC/Rcok Creek