SOLUTIONS UNIT
Name:_________________________________________________Period:____Date:____
I.
MIXTURES:
1) ____________________ = a blend of two or more kinds of matter, each of which retains its
own identity and properties
 can be physically separated (filtration, evaporation, decanting, magnetism, etc)
a) ___________________________________= a mixture that is uniform in composition
throughout
 Ex: Food coloring and water
b) ___________________________________= a mixture that is NOT uniform in
composition throughout
 Ex: Oil and water
A. Types of Mixtures:
1) ______________________________= a homogeneous mixture
2) ______________________________= a mixture in which the particles are so large that they
settle out unless the mixture is constantly stirred or agitated
 Heterogeneous mixture
 Ex: Sand and water
3) ____________________________= a mixture consisting of particles that are intermediate in
size between those in solutions and those in suspensions
 Heterogeneous mixture
 Ex: Milk, mayonnaise, smog, butter, whipped cream
1
a.) What property of a colloid helps to prevent colloid particles from settling out of a mixture?
 ___________________________________= the random continuous motions of
colloidal particles
b.) _____________________________________= visible pattern caused by the reflection
of light from suspended particles in a colloid (or from suspended particles in a suspension if
the particles have not settled out)

Ex: visibility of a headlight beam on a foggy night
B. Comparison of solution, colloids, and suspensions
Solutions
Colloids
Suspensions
Heterogeneous
Heterogeneous
Particle size: 1-1000 nm,
Particle size: over 1000 nm,
dispersed; can be aggregates or
suspended; can be large particles
large molecules
or aggregates
_______ separate on standing
________ separate on standing
Particles settle out
Cannot be separated by
Cannot be separated by
filtration
filtration
___________ scatter light
Scatter light (Tyndall effect)
_________________
Particle size: 0.01-1 nm; can be
atoms, ions, molecules
_____ be separated by filtration
______ scatter light, but are not
transparent
C. Determining if a mixture is a true solution, a colloid, or a suspension:
1) If particles settle or can be filtered out = ___________________
2) If particles DO NOT settle or filter out shine a beam of light (Tyndall effect) through the
mixture
 If the Tyndall effect is observed = ____________________________

2)
If the Tyndall effect is NOT observed = _________________________
THE NATURE OF SOLUTIONS:
1) ______________ = the substance that does the dissolving in a solution
a) Typically present in the greatest amount
b) Typically a liquid
c) ____________________ is the most common or “universal” solvent

B/c water molecules are __________________________
2

The hydrogen side of each water (H2O) molecule carries a slight positive electric
charge, while the oxygen side carries a slight negative electric charge.

Water can dissociate ionic compounds into their positive and negative ions

The positive part of an ionic compound is attracted to the oxygen side of water while
the negative portion of the compound is attracted to the hydrogen side of water.

Water won't dissolve or won't dissolve well. If the attraction is high between the
opposite-charged ions in a compound, then the solubility will be ________________.
o
Ex: hydroxides exhibit low solubility in water.
o
Ex: nonpolar molecules don't dissolve very well in water (fats and waxes)
2) ________________________ = substance being dissolved in a solution
a) Typically present in the least amount
b) Typically a solid
A.9 Possible Solution Combinations:
Solvent
Gas
Solute
Common Example
Gas
Diver’s tank
Gas
Liquid
Gas
Solid
Liquid
Gas
Liquid
Liquid
Liquid
Solid
Solid
Gas
Solid
Liquid
Solid
Solid
Moth ball
Vinegar
o NOT all solutions are
liquids/solids!
o Solutions are formed
in ALL 3 states!
Gas stove lighter
Sterling Silver (Ag + Cu)
C. Solvation:
1) ___________________________________ = the process of dissolving
a.) First- solute particles are surrounded by solvent particles
b.) Then- solute particles are separated and pulled into solution
2) ____________________________________= separation of an ionic solid into aqueous ions
3



Ex: NaCl + H2O – the Na ion and Cl ion become hydrated and gradually move away from
the crystal into solution.
Each ion in the solution acts as though it were present________________: So there is
only a solution containing Na+ and Cl- ions uniformly mixed with H2O particles
NaCl(s)  Na+(aq) + Cl–(aq)
3) ____________________________= breaking apart of some polar molecules into aqueous ions

Ex: HNO3(aq) + H2O(l)  H3O+(aq) + NO3–(aq)
4) _________________________________________ =molecules stay intact

Ex: C6H12O6(s)  C6H12O6(aq)
D. Factors Affecting the Rate of Dissolving (Increases Solution Rate):
1) ____________________: increases surface area exposed to solvent
2) ____________________: allows solvent continual contact with solute
3) ____________________: increases kinetic energy; increases mixing
3)
Electrolytes and Nonelectrolytes
A. Electrolytes and Nonelectrolytes
1)
________________________ = a substance that dissolves in water to give a solution that
conducts electric current
2) _________________________ = a substance that dissolves in water to give a solution that
does NOT conduct an electric current
3) Solutions of electrolytes can conduct electric current:
a) The positive ions and the negative ions ___________________________(separate) in
solution. The ___________________ions can move a charge from one point in the solution
to another point
4) Solutions of nonelectrolytes CANNOT conduct electric current:
a) When a nonelectrolyte dissolves in water there are NO
_________________________________ in solution.
b) Ex: Solute exists as molecules
5) Weak Electrolytes
a) Only a portion of dissolved molecules ionize.
6) Solid ___________________ compounds cannot conduct electric current:
a) Ions are present but they are NOT ___________________.
4
4)
SOLUBILITY:
1) _________________________= quantity of solute that will dissolve in specific amount of solvent
at a certain temperature. (Pressure must also be specified for gases).
a) Ex: 204 g of sugar will dissolve in 100 g of water at 20C
a) soluble and insoluble are relative terms
b) solubility should NOT be confused with the rate at which a substance dissolves
2) _______________________ = a stable solution in which the maximum amount of
solute has been dissolved.
a) Visual evidence: a quantity of undissolved solute remains in contact with
the solution
3) __________________________________ = state where the solute is
dissolving at the same rate that the solute is coming out of solution (crystallizing).
a) Opposing processes of the dissolving and crystallizing of a solute occur at equal rates.
b) solute + solvent
solution
4) _________________________________ = a solution that contains less solute than a saturated
solution under existing conditions
5) __________________________________= a solution that temporarily contains more than the
saturation amount of solute than the solvent can hold.
 Unstable – if disturbed, the excess solute will crystallize out of solution
B. 3 FACTORS EFFECTING SOLUBILITY:
1) Nature of solute and solvent
a) “____________________________” = rule of thumb for predicting whether or not one
substance dissolves in another

“Alikeness” depends on:
o Intermolecular forces
o Type of bonding
o Polarity or nonpolarity of molecules:
 ________________ solutes tend to dissolve in polar solvents but not in nonpolar solvents
b) Solvent-Solute Combinations:
Solvent Type
Solute Type
Polar
Polar
Polar
Nonpolar
Nonpolar
Polar
Nonpolar
Nonpolar
Is Solution Likely?
“Like Dissolves Like”
5
2) __________________________:
a) Pressure has __________________ on the solubility of liquids or solids in
liquid solvents.
b) The solubility of a _________ in a ______________ INCREASES when
pressure __________________. It is a direct relationship.
3) ______________________________:
a) The solubility of a gas in a liquid solvent DECREASES with an ______________ in
temperature.
b) The solubility of a solid in a liquid solvent MOST OFTEN increases with an increase in
temperature. However, solubility changes vary widely with temperature changes sometimes
decreasing with temperature increases.
5)
SOLUBILITY GRAPH:
A. Solubility Curve
1) ______________________ = any point on the line or ABOVE the line
2) ___________________________ = any point
BELOW the line
B. Solubility Curve Problems
1) What is the solubility of the following solutes in
water?
a)
NaCl at 60ºC = ____________
b)
KCl at 40ºC = _____________
c)
KNO3 at 20ºC = ____________
2) Are the following solutions saturated or
unsaturated? Each solution contains 100 g of H20.
a)
31.2 g of KCl at 30ºC =
_______________
b)
106g KNO3 at 60ºC = _______________
c)
40 g NaCl at 10ºC = _______________
d)
150 g KNO3 at 90ºC = _______________
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3) For each of the following solutions, explain how much of the solute will dissolve and how much will
remain undissolved at the bottom of the test tube?
a) 180 g of KNO3 in 100 g of water at 80ºC
______________________________
______________________________
b) 180 g of KNO3 in 100 g of water at 20ºC
_______________________________
_______________________________
c) 60 g of NaCl in 100 g of water at 60ºC
________________________________
________________________________
4) A saturated solution of KNO3 is formed from one hundred grams of water. If the saturated solution
is cooled from 90°C to 30°C, how many grams of precipitate are formed?
_______________________________
5) A saturated solution of KCl is formed from one hundred grams of water. If the saturated solution is
cooled from 90°C to 40°C, how many grams of precipitate are formed?
__________________________________
6)
CONCENTRATION: MOLARITY & MOLALITY
1) ______________________________= The amount of solute in a solution.
a) Describing Concentration:
• % by mass - medicated creams
• % by volume - rubbing alcohol
• ppm, ppb - water contaminants
• molarity - used by chemists
• molality - used by chemists
A. Molarity:
1) ___________________________= unit of concentration of a solution expressed in moles of
solute per liters of solution.
• Formula:
M = ___MOLES of solute______
LITERS of solution
 Ex: 2M HCl
o
What does that mean?
Mol

M=

2M HCl =
L
2 mol HCl
1L
7
B. Molarity Calculations:
1) To solve for Molarity when given grams of solute, you will use gram/mole conversion and the
equation for Molarity. Use the following steps:
a)
Change grams of solute to moles. To do this you need to use the following conversion:
Grams of solute
1 mole of solute
=
moles of solute
Molar mass of solute*
*make sure formulas are correct; Use ion list!
b) Change the volume of solution to liters (L) by moving the decimal 3 places to the LEFT.

Remember that 1000 mL = 1 L. Water will most often be the solvent in a solution.
c) Substitute the information into the molarity equation and solve.
2) Ex: What is the molarity of a solution composed of 22.4g of sodium chloride dissolved in
enough water to make 500mL of solution?
22.4 g of NaCl
1 mole_NaCl
=
0.383 moles NaCl
58.443 g NaCl
0.383 moles NaCl = 0.766 M NaCl
Note: You need to change mL of solution to L. So
move decimal 3 places to the left!
0.500 L solution
3) Use the Molarity equation to solve for grams of solute. You will need to solve the Molarity
equation for moles of solute. Then convert moles of solute to grams of solute.
4) Ex: How many grams of lithium bromide are present in 300mL of a 0.4M lithium bromide
solution?
a) Substitute the information and solve for moles by cross multiplying or multiplying by the
reciprocal:
“X” moles LiBr___
= 0.4M LiBr
0.300 L solution
X = 0.12 moles of LiBr are needed
b) Convert moles to grams using the following conversion:
Moles of solute
molar mass of solute
= grams of solute
1 mole of solute

The substitution would look like:
0.12 moles LiBr
86.845 g LiBr
= 10. g LiBr
1 mol LiBr
8
C. Dilution Calculations:
1) Sometimes solutions of lower concentrations are made from existing solutions.
2) Moles of solute remain the same.
3) Formula: M1V1 = M2V2
 M1 is the original molarity concentration
 V1 is the volume of the original solution (in L)
 M2 is the new concentration
 V2 is the amount of the new solution needed (in L)
4) Ex 1: What volume of 15.8M HNO3 is required to make 250 mL of a 6.0M solution?
GIVEN:
WORK: M1 V1 = M2 V2
M1 = 15.8M
V1 = ?
(15.8M) V1 = (6.0M)(250mL)
M2 = 6.0M
V2 = 250 mL
V1 = 95 mL of 15.8M HNO3
D. Molality:
1) _____________________________(m) = a unit of concentration of a solution expressed in
moles of solute per kilogram of solvent.
a) Ex: A solution that contains 1 mol of solute, ammonia (NH3), dissolved in
exactly 1 kg solvent, is a “one molal” solution.
b) Formula:
m = _MOLES of solute
kg of solvent
c) Note: 1 kg water = 1 L water
E. Calculations: Solving for Molality:
1) If needed convert:

grams of solute  moles of solute

mL of solvent  kg of solvent (move decimal 3 places to the left!)
2) Solve for molality
m = moles of solute
kg of solvent
3) Ex: Find the molality of a solution containing 75 g of MgCl2 in 250 mL of water.
9
F. Calculations: Using the Molality equation to solve for grams of solute:
1)
Use Molality equation and solve for moles
2)
Then convert moles of solute to grams of solute
3)
Ex: How many grams of NaCl are required to make a 1.54m solution using 0.500 kg of water?
COLLIGATIVE PROPERTIES:
A. Colligative Property:
1)
___________________________________= A property that depends on the NUMBER of
solute particles rather than the type of particle.
B. Examples of colligative properties:
1) ______________________________________ LOWERING of a solution
a) Ex: We have two beakers. One contains H2O, the other contains a salt solution (NaCl) in
H2O. Both beakers are placed into a sealed chamber:
DAY 3: If we leave it
for a couple of days
and then come back
and take a look, this is
what it might look like.
DAY 1
b) WHAT HAPPENED?:



Na+ and Cl- ions are non-volatile ions. In other words, they will not leave and go into
the vapor phase.
Those H2O molecules with enough kinetic energy will leave the surface of the solutions
and enter into the vapor phase.
The fact that Na+ and Cl- ions are dissolved in H2O indicates that there is an
attractive interaction between the solutes and the H2O.
10





The interaction of the Na+ and Cl- ions for H2O will act to hinder the ability of the
solvated H2O molecules to leave and go into the vapor phase (Na+ and Cl- ions are nonvolatile).
This hindrance of the H2O molecules to enter the vapor phase will reduce the vapor
pressure of the H2O in the salt-containing solution.
The two beakers are in the same sealed container, thus, the vapor pressure above the
solutions is identical.
The rate of H2O entering the solutions by collisions from the vapor state will be
identical
The rate of H2O leaving the liquid phase and entering the vapor phase is slower for
the NaCl-containing solution
2) _______________________________________ = If two solutions, with different solute
concentrations, are separated by a semi-permeable membrane, there can be a net flow of
solvent across the membrane.
3) _________________________________________= The solution will begin to freeze at a
temperature BELOW that of the pure solvent.
4) _________________________________________ = the solution will begin to boil at a
temperature ABOVE that of the pure solvent.
C. How are solutions different than pure liquids?
1) One of the ways in which they are different, is that when you add a solute to a liquid _________
the freezing point and boiling point of the solution ________________________.
a) ____________________________________ – In a solution, the solute particles
interfere with the attractive forces among the solvent particles.

This prevents the solvent from entering the solid state at its normal freezing point
b) ____________________________________- When the temperature of a solution
containing a nonvolatile solute is raised to the boiling point of the pure solvent, the
resulting vapor pressure is still less than the atmospheric pressure and the solution will
not boil.

Thus, the solution must be heated to a higher temperature to supply the additional
KE needed to raise the vapor pressure to atmospheric pressure.
2) Water is the liquid we will be dealing with most often
a) The freezing point of pure water is _______.
b) The normal boiling point of water _________.
c) But if you make a solution using water as the _____________________, the freezing
point of that solution will _________ be 0°C nor will the boiling point be 100°C!
D. Freezing Point Depression & Boiling Point Elevation:
1) ____________________________________________(TFP) – solutions will
freeze at LOWER temperature than the pure solvent
a) The more solute dissolved, the _________________the effect.
11
b) Ex: ethylene glycol (antifreeze) protects against freezing of the water in the cooling
system by lowering the freezing point to about -40°C
c) Ex: making of homemade ice cream- The ice cream mix is put into a metal container which
is surrounded by crushed ice. Then salt is put on the ice to lower its melting point. This
gives a temperature gradient across the metal container into the saltwater-ice solution
which is lower than 0°C. The heat transfer out of the ice cream mix allows it to freeze.
2) _______________________________________________(TBP) – solutions will boil at
HIGHER temperatures than the pure solvent
a) The boiling point of pure water is 100°C, but that boiling point can be elevated by the
addition of a solute such as a salt.
b) The more solute dissolved, the _____________________the effect.
E. Calculating Freezing Point Depression & Boiling Point Elevation:
1) Solution concentrations are given in molality (m)
2) Colligative properties are ______________ proportional to the molal concentration of a solute
3) Account for particle molality for ELECTROLYTES
 Ex: NaCl = 2 ions (Na+ & Cl-)
 Ex: MgCl2 = 3 ions (Mg+2 and 2 Cl-)
4) A change in the concentration (m), changes the freezing point and boiling point of a solution
a) You will need to figure out the change in temperature for the normal freezing point (FP) or
boiling point (BP) and adjust them with the following steps:
∆T = mK
∆T = change in temperature
m = moles solute/kg solvent (MOLALITY)
K = constant
 You will mostly use the constants for water:
Kfp = 1.853 ˚C/m
Kbp = 0.515˚C/m
 Other values can be found below:
FREEZING Point Depression Constants
BOILING Point Elevation Constants
Compound
Freezing Point (oC)
Kfp (oC/m)
Compound
Boiling Point (oC) Kbp(oC/m)
water
0.000
1.853
water
100.000
0.515
acetic acid
16.660
3.90
ethyl ether
34.550
1.824
benzene
5.530
5.12
carbon disulfide
46.230
2.35
p-xylene
13.260
4.3
benzene
80.100
2.53
naphthalene
80.290
6.94
carbon tetrachloride
76.750
4.48
cyclohexane
6.540
20.0
camphor
207.420
5.611
12
4) Steps solving for ∆T:
a) Solve for molality (see steps from previous problems)
m = moles of solute
kg of solvent
b) Find your K constant from the chart listed in your notes.

Be sure to find the correct K constant for what you are solving for-- either FP or BP!!
c) Solve for ∆T
∆T = mK
d) If solving for:
 Freezing point: Take the normal freezing point of the solvent and SUBTRACT the ∆T value
o Round final answer to 3 significant figures!
 Boiling point: Take the normal boiling point of the solvent and ADD the ∆T value
o Round final answer to 6 significant figures!
5) Ex #1: At what temperature will a solution that is composed of 32.8 g of glucose (C6H12O6) in 225 g of
water boil?
6) Ex #2: Find the freezing point of a saturated solution of NaCl containing 28.0 g NaCl in 115 g water.
13
Solubility Curve Worksheet
1) Look at the graph below. In general, how does temperature affect solubility?
________________________________________________________________________________
2) Which compound is LEAST soluble at 10 °C? ______________________
3) How many grams of KCl can be dissolved in 100g of water at 80°C? ______________________
4) How many grams of NaCl can be dissolved in 100g of water at 90°C? ______________________
5) At 40°C, how much KNO3 can be dissolved in 100g of water? ___________
6) Which compound shows the least amount of change in solubility from 0°C-100°C? ________________
7) At 30°C, 90g of NaNO3 is dissolved in 100g of water. Is this solution saturated or unsaturated?
_____________
8) At 60°C, 72g of NH4Cl is dissolved in 100g of water. Is this solution saturated or unsaturated?
_____________
9) A saturated solution of KClO3 is formed from one hundred grams of water. If the saturated solution is
cooled from 90°C to 50°C, how many grams of
precipitate are formed?_______________
10) A saturated solution of NH4Cl is formed from one
hundred grams of water. If the saturated solution
is cooled from 80°C to 40°C, how many grams of
precipitate are formed?_______________
11) Which compounds show a decrease in solubility
from 0°C-100°C?
_____________________________
12) Which compound is the most soluble at 10°C?
_____________________________
13) Which compound (besides Ce2(SO4)3) is the least
soluble at 50°C?________________________
14) For each of the following solutions, explain how
much of the solute will dissolve and how much will
remain undissolved at the bottom of the test tube?
a) 120 g of KCl in 100 g of water at 80ºC
_______________________________
_______________________________
b) 130 g of NaNO3 in 100 g of water at 50ºC
_______________________________
_______________________________
14
Molarity Worksheet
1) What is the molarity of a solution that is composed of the following?:
a) 30g of calcium chlorate dissolved in enough water to make 370mL of solution
b) 450g of potassium iodide dissolved in enough water to make 4L of solution
c) 25g of hydrochloric acid dissolved in enough water to make 600. mL of solution
15
d) 0.750g of aluminum sulfate dissolved in enough water to make 75.0 mL of solution
e) 35.5g of silver nitrate dissolved in enough water to make 125mL of solution
2) How many grams of each of the following solute are needed to make the solution:
a) 215 mLof 0.400M iron III nitrate
b) 565 mL of 3.00M potassium iodide
16
c) 355 mL of 0.250M chromium II permanganate
d) 400.0 mL of 6.00M zinc sulfate
e) 800.0 mL of 1.00M magnesium acetate
17
Dilutions Worksheet
1) What formula is used when solutions of lower concentrations are made from existing solutions?
2) What volume of 5.30M ammonium phosphate is required to make 200.mL of a 4.25M solution?
3) What volume of 2.50M mercury I chromate is required to make 225mL of a 0.200M solution?
4) What volume of 0.5M copper II cyanide is required to make 25mL of a 0.10M solution?
5) What volume of 1.50M tin II fluoride is required to make 250. mL of a 0.875M solution?
6) What volume of 2.55M manganese dichromate is required to make 750. mL of a 0.150M solution?
18
Molality (m) Worksheet
1) Determine the molality of each of the following solutions:
a) 199g of nickel II bromide in 500. g H2O
b) 4.5 mol Mg(NO3)2 in 2.5 kg H2O
c) 0.625 mol K2SO4 in 850. g H2O
d) 92.3g of potassium fluoride in 137g H2O
19
e) 85.2g tin II bromide in 142g H2O
2) How many grams of KCl are required to make a 0.525m solution using 250. g of water?
3) How many moles of each solute would be required to prepare each of the following solutions?
a) 0.45 m CaSO4 in 2.5 kg H2O
b) 3.25 m Ba(NO3)2 in 750. g H2O
20
4) Determine how many grams of solute is required to make each of the following solutions:
a) A 4.50 m solution of H2SO4 in 1.25 kg H2O
b) A 3.50 m solution of MgCl2 in 0.450 kg H2O
5) A solution is prepared by dissolving 17.1 g of sucrose (C12H22O11), in 275 g of H2O. What is the molality of that
solution?
21
6) How many kilograms of H2O must be added to 75.5 g of Ca(NO3)2 to form a 0.500 m solution?
7) How many grams of glucose (C6H12O6), must be added to 750. g of H2O to make a 1.25 m solution?
22
Freezing Point Depression & Boiling Point Elevation Worksheet:
Directions: Round final answer to 3 significant figures for freezing point depression problems and round final
answer to 6 significant figures for boiling point elevation!
1) If 68.4g of sucrose C12H22O11 are dissolved in 500. g of water, what will be the boiling and freezing points of the
solution?
2) Calculate the freezing point of a solution containing 5.75 g of sugar C12H22O11 in 59.5 g of water.
3) Calculate the freezing point of solution that contains 32.7 g of sodium acetate dissolved in 250. g of water.
23
4) Calculate the freezing point, if 46.5 g of glycerol, C3H5(OH)3, are dissolved in 500. g of water.
5) Calculate the boiling point of 30.8 g of silver nitrate in 500. g water.
6) Determine the boiling point of 16.3 g of calcium chloride dissolved in 250. g water.
7) Determine the boiling point of 23.9 g of copper (II) sulfate in 300. g water
24
Review Molarity, Dilutions, Molality, Freezing Point Depression & Boiling Point Elevation:
1) What is the molarity of a solution that is composed of the following?:
a) 26g of calcium chloride dissolved in enough water to make 425mL of solution.
b) 329g of calcium iodide dissolved in enough water to make 2.5L of solution.
2) How many grams of each of the following solute are needed to make the following solutions?:
a) 500. mL of 3.40M iron (III) nitrate.
b) 700. mL of 0.880M lithium sulfate
25
3) What volume of 6.25M ammonium chloride is required to make 500. mL of a 3.92M solution?
4) What volume of 5.50M sodium dichromate is required to make 725mL of a 0.600M solution?
5) Determine the molality of each of the following solutions:
a) 254 g of NaCl in 575 g water
b) 9.65 mol MgCl2 in 675 kg water
26
6) Determine the freezing point and boiling point of a solution containing 98.5 g of NaCl in 100. g water.
7) Determine the freezing point and boiling point of a solution containing 425 g of CaCl 2 in 550. g water.
27
SOLUTIONS CONCEPT REVIEW QUESTIONS:
1) Explain why solutions are classified as mixtures instead of compounds.
2) What are the three different types of mixtures?
3) Differentiate between a homogeneous mixture and a heterogeneous mixture.
4) Classify each of the following as a heterogeneous mixture or a homogeneous mixture.
a) Salad __________________________________
b) Tap water _______________________________
c) Muddy water ____________________________
5) What another name for a solution? ______________________________
6) What is the difference between a solute and a solvent?
7) What is the Tyndall Effect? Give a common example of this effect.
8) In what type of mixture is to easiest to separate the component substance? Why?
9) How does a solution behave differently from a suspension when a beam of light is shined through it?
28
10) What property of a colloid helps to prevent colloid particles from settling out of a mixture?
11) Given an unknown mixture consisting of two substances, explain how a scientist could use lab
techniques to determine whether the mixture is a true solution, a colloid, or a suspension.
12) What is considered to be the “universal” solvent? ___________________
a) Explain why.
13) What is the solute in a brass alloy containing 75% copper and 25% zinc? ____________________
a) How do you know?
14) Not all solutions are solids dissolved in liquids. Provide two examples of other types of solutions.
15) Suppose you were making a salt water solution. Sodium ions and chloride ions are separating from one
another.
a) What is this called? _____________________________________________
16) Suppose you were making a salt water solution. Sodium ions and chloride ions are each being
surrounded by water molecules.
b) What is this called? _______________________________________
17) Describe how stirring increases the rate of dissolving of a solute.
18) Describe how heating increases the rate of dissolving of a solute.
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19) Describe how grinding a solute increases the rate of dissolving of a solute.
20) Define solubility.
21) What are 3 factors that affect solubility?
a) ________________________________________
b) ________________________________________
c) ________________________________________
22) What happens to the solubility of a gas in a liquid solvent as temperature increase?
23) Explain the rule, “Like Dissolves Like”.
24) State whether each of the following will conduct an electric current. Also, explain why each does or
does not conduct an electric current.
a) salt (NaCl) water
b) sugar (C12H22O11) water
c) solid NaCl
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25) When does solubility equilibrium occur?
26) What are the differences between a saturated solution, unsaturated solution and a supersaturated
solution?
27) How could you tell by looking at a solution that it was saturated?
28) Differentiate between molarity and molality.
29) What does 5M HNO3 mean?
30) What is the molarity formula?
31) What is the molality formula?
32) What formula is used when solutions of lower concentrations are made from existing solutions?
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33) Define colligative property.
34) Provide 4 examples of colligative properties.
35) Explain how the boiling point and freezing point of solutions are different than pure liquids?
 Use the solubility cure below to answer the following questions:
36) Which salt is LEAST soluble at 20 °C? ______________________
37) How many grams of KBr can be dissolved in 100g of water at 60°C? ______________________
38) How many grams of NaCl can be dissolved in 100g of water at 100°C? ______________________
39) At 40°C, 180g of NaClO3 is dissolved in 100g of water. Is this solution saturated or unsaturated?
_____________
40) At 70°C, 70g of KBr is dissolved in 100g of
water. Is this solution saturated or
unsaturated? _____________
41) A saturated solution of NaClO3 is formed
from one hundred grams of water. If the
saturated solution is cooled from 80°C to
60°C, how many grams of precipitate are
formed?_______________
42) How much of the solute will dissolve and
how much will remain undissolved at the
bottom of the test tube?
d) 160 g of KNO3 in 100 g of water at
50ºC
____________________________
____________________________
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SOLUTIONS TEST STUDY GUIDE
Chapter 13 (pg 513-547)
 Read over your notes and rework your class/homework assignments. Study your lab reports.
You will do “Mole-tastic” on this test if you can do the following things.
Mixtures
 Explain the difference between mixtures and compounds?
 Explain the difference between mixtures (solutions, colloids, suspensions).
 What can be done experimentally to differentiate between the 3 types of mixtures?
 Examples of each type of mixture
 Define the Tyndall Effect.
Nature of Solutions




Explain the difference between solute and solvent?
Explain why is water considered the “universal solvent”?
Describe at a molecular level the dissolving of an ionic compound in water.
Explain the 3 factors that can increase the solution rate.
Electrolytes & Non-electrolytes
 Differentiate between electrolytes and non-electrolytes.
 What properties do electrolytes have that are not present in non-electrolytes.
Solubility




Define solubility.
Explain the difference between a saturated, unsaturated and a supersaturated solution.
Explain the general rule of thumb, “like dissolves like”.
Explain the 3 factors that affect solubility.
Solubility Graph
 Be able to interpret a solubility curve
Concentrations
 Explain the difference between Molarity and Molality
 Be able to solve Molarity problems
 For Molarity
 For grams of solute
 Be able to calculate how to dilute a solution
 Be able to solve Molality problems
 For Molality
 For moles of solute
 For grams of solute
Colligative Properties
 Define colligative property.
 Describe examples of colligative properties.
 How are solutions different than pure liquids?
 Explain what happens to boiling and freezing points
 Calculate freezing point and boiling point of solutions.
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