Name:
Date:
Unit 9 – Liquids and Solutions
Accelerated Chemistry I
Phase Changes
1. Name the phase change in each of these events:
a. Dew appears on a lawn in the morning.
b. Icicles change into liquid water.
c. Wet clothes dry on a summer day.
2. Name the phase change in each of these events:
a. A diamond film forms on a surface from gaseous carbon atoms in a vacuum.
b. Mothballs in a bureau drawer disappear over time.
c. Molten iron from a blast furnace is cast into ingots (“pigs”).
3. a. Why are gases more easily compressed than liquids?
b. Why do liquids have a greater ability to flow than solids?
4. Liquid propane, a widely used fuel, is produced by compressing gaseous propane at 20°C. During the
process, approximately 15 kJ of energy is released for each mole of gas liquefied. Is this process
endothermic or exothermic? Where does this energy come from?
Vapor Pressure
5. Compare the evaporation of a liquid in a closed container with that of liquid in an open container.
6. Describe what is happening at the molecular level when a dynamic equilibrium occurs.
7. Explain why increasing the temperature of a liquid increases its rate of evaporation.
8. Distinguish between the boiling point and the normal boiling point of a liquid.
9. When you remove the lid from a food container that has been left in a freezer for several months, you
discover a large collection of ice crystals on the underside of the lid. Explain what has happened.
10. Explain why a liquid stays at a constant temperature while it is boiling.
Accel Unit 9 Student Handout
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11. Refer to the following figure to answer these questions about chloroform, ethanoic acid, water, and ethanol.
a. What is the normal boiling point of ethanoic
acid?
b. Which liquid has the highest vapor pressure
at 40°C?
c. Which of the liquids would be easiest to
evaporate?
d. At standard atmospheric pressure, which of
the substances are in the gaseous state at
70°C?
e. Water boils at l00°C at standard pressure. How would the pressure on ethanol and on ethanoic acid have
to change for these liquids to boil at l00°C?
f. Which liquid experiences the strongest intermolecular forces?
12. Why are pressure cookers recommended for cooking at high altitude?
13. The table gives the vapor pressure of isopropyl alcohol at various temperatures. Graph the data. Use a
smooth curve to connect the data points.
Temperature, °C Vapor Pressure, mm Hg
0
8.33
25
45.2
50
179
75
565
100
1480
125
3390
a. What is the estimated normal boiling point of
isopropyl alcohol?
b. What is the boiling point of isopropyl alcohol
when the external pressure is increased to twice
standard pressure?
c. What is the pressure when isopropyl alcohol boils
at 60°C?
d. The chart shows the vapor pressure of isopropyl alcohol at temperatures above the normal boiling point.
Describe the physical state of a sample of isopropyl alcohol at 125°C, when the vapor pressure is 3390
mm Hg?
Accel Unit 9 Student Handout
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G
b. freezing begins
c. boiling begins
d. condensation begins
Temperature
Use the following figure to answer the next 3 questions.
E
14. Identify the point on the graph where each of the
B following occurs.Figure 2.
a. melting begins
A
C
D
F
Gas
Liquid→Gas
Liquid
Solid
15. Does the substance give off or take in heat as it goes
from D to C?
Solid→Liquid
Energy
16. Explain why the temperature remains constant
during a phase change from C to D, even though heat
flows from the surroundings to the system.
17. Evaluate this statement: the energy content of a substance is higher in the liquid phase than in the vapor
phase at the same temperature.
Heat of Fusion/Vaporization
Use the graph above (Figure 2.)and table to the right
for the next 4 questions. Assume 1.00 mol of substance
in each container.
18. Calculate heat absorbed in region C→ D for neon.
Answ: 0.33 kJ
Ammonia
Benzene
Methanol
Neon
Water
Freezing
point
(K)
195.3
278.7
175.5
24.5
273.0
ΔHfus
(kJ/mol)
5.65
9.87
3.16
0.33
6.02
Boiling
point
(K)
239.7
353.3
337.2
27.1
373.0
19. Calculate heat absorbed in region E→ F for benzene. Answ: 30.8 kJ
20. Calculate heat absorbed in regions D→ F for methanol (specific heat = 81.6 J/g⋅°C). Answ: 458.2 kJ
21. Calculate heat absorbed in regions C→ F for ammonia (specific heat = 35.1 J/g⋅°C). Answ: 55.7 kJ
Accel Unit 9 Student Handout
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ΔHvap
(kJ/mol)
23.4
30.8
35.3
1.76
40.7
Pressure (not to scale)
d
PC
e

b
f
i
22. How many grams ice would ahave
 to melt to lower the temperature of 352 g of water from 25°C to 5°C?
Answ: 88 g
g
c
h
Temperature (not to scale)
TC
23. How many kJ of heat is required to warm 10.0 g of ice, initially at -10°C, to steam at 110.0°C. The heat
capacity of ice is 2.09J/g⋅°C and that of steam is 2.01 J/g⋅°C. (Use your notes for this question!) Answ:
3.05x104 J = 30.5 kJ
P HASE D IAGRAMS
24. Consider the phase diagram for substance X shown below. Identify the phases present at points a through g.
a.
e.
b.
f.
c.
g.
d.
h. Describe what happens when you start at point a and increase
the temperature at constant pressure.
i. Describe what happens when you start at point b and decrease
the pressure at constant temperature.
j. Is liquid X more or less dense than solid X?
Accel Unit 9 Student Handout
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Iodine (I2)
b. What is the melting point for iodine at 1 atm?
Pressure (not to scale)
PC
25. Consider the phase diagram for iodine shown to the right and
answer each of the following questions.
a. What is the normal boiling point for iodine?
1 atm
c. What phase is present at room temperature and normal
atmospheric pressure?
d. What phase transitions occur as you uniformly increase the
pressure on a gaseous sample of iodine from 0.010 atm at
185°C to 100 atm at 185°C?
113.6°
184.4°
TC
C
Temperature,°C C(not to scale)
C
C
26. Nitrogen has a normal boiling point of 77.3K and a melting point (at 1 atm) of 63.1K. Its critical
temperature is 126.2K and critical pressure is 2.55104 mm Hg. It has a triple point at 63.K and 94.0 mm
Hg. Sketch the phase diagram for nitrogen. Does nitrogen have a stable liquid phase at 1 atm?
Solutions
1. For each of the following solutions, identify which substance is the solute and which is the solvent.
a. 20 g NaCl and 100 g H2O
b. 50 g benzene and 1.0 g aspirin
c. 100 g H2 and 50 g O2
d. 25 g CH3COCH3 and 100 g H2O
2. Explain the difference between a strong electrolyte solution and a non-electrolyte solution. What kinds of
solutes form strong electrolyte solutions?
Unit 9 Student Handout
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3. What is solvation? What are the factors that influence the extent to which solvation occurs? Give two
examples of solvation, including one that involves ion-dipole interaction and another in which dispersion
forces come into play.
4. You are given a bottle of solid X and three aqueous solutions of X—one saturated, one unsaturated, and one
supersaturated. How would you determine which solution is which?
5. How is rock candy made?
6. Explain the variations in solubility in water of the alcohols listed here:
Compound
Solubility in Water (g/100 g, 20°C)
CH3OH
∞ [means completely miscible]
CH3CH2OH
∞
CH3CH2CH2OH
∞
CH3CH2CH2CH2OH
9
CH3CH2CH2CH2CH2OH
2.7
7. Choose the member of each set that you would expect to be more soluble in water. Explain your answer.
(Hint: Which molecule has IMF MOST similar to that of water?)
a. chloromethane, CH3Cl, or methanol, CH3OH
b. nitrogen triiodide or potassium iodide
c. lithium chloride or ethyl chloride, C2H5Cl
d. ammonia or methane
e. naphthalene, C10H8, or hydrogen peroxide, H-O-O-H
f. methyl alcohol, CH3OH, or dimethyl ether, H3C-O-CH3
Unit 9 Student Handout
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

8. One mole of CaCl2 is represented as
where  represents Ca and  represents Cl. Complete the
picture
showing
only
the
calcium
and
chloride
ions. The water molecules need not be shown. What is the



 2+
1.00
molarity
of Ca ?+of
Cl-L
?
H 2O





Net Ionic Equations.
Equations which involve ions and exclude any species that do not take part in
the reaction are referred to as net ionic equations. Like all equations, net ionic equations must show atom
balance and charge balance (same total charge on both sides).
Strong Acids – Present in solution largely as ions rather than as molecules. The six strong acids and their
formulas must be memorized.
hydrochloric - HCl
hydrobromic - HBr
hydroiodic - HI
nitric - HNO3
sulfuric - H2SO4
perchloric - HClO4
Strong Bases – Present in solution largely as hydroxide ions and metal ions. Memorize the six strong bases and
their formulas.
lithium hydroxide - LiOH
sodium hydroxide - NaOH
potassium hydroxide – KOH
calcium hydroxide - Ca(OH)2
strontium hydroxide - Sr(OH)2
barium hydroxide - Ba(OH)2
In the following reactions, first predict the products (not all reactions “go”). Then
write a balanced net ionic equation for the reaction, if it applies. Remember to
indicate the physical state of each of the reactants and products.
1. Solid lithium reacts with nitrogen gas to produce a solid.
2. Aqueous potassium chlorate decomposes to aqueous potassium chloride and a gas.
3. Liquid water plus carbon dioxide gas in a synthesis reaction yielding a weak acid present in seltzer water.
4. Liquid octane combusts.
Unit 9 Student Handout
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5. Liquid 2-methyl-nonane combusts.
6. Solid zinc is dropped into aqueous hydrochloric acid.
7. Liquid bromine is added to aqueous potassium iodide.
8. A piece of aluminum is dropped into aqueous acetic acid.
9. Iron filings are mixed with aqueous lead(IV) nitrate to form an iron(II) product.
10. Solid potassium is dropped into water.
11. Solid copper is mixed with aqueous aluminum sulfate.
12. Aqueous zinc nitrate is mixed with aqueous potassium carbonate.
13. Aqueous sodium chloride is mixed with aqueous sulfuric acid.
14. Solid lithium oxide reacts with aqueous hydrofluoric acid in a double replacement reaction.
Unit 9 Student Handout
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15. Solid silver is mixed with zinc chloride solution.
G ASES
DISSOLVED IN
S OLUTION
9. Explain why warm soda goes flat faster than cold soda.
10. Explain why a can of soda fizzes when opened.
11. For a saturated aqueous solution of each of the following at 20° C and 1 atm, will the solubility increase,
decrease, or stay the same when the indicated change occurs?
a. O2(g), increase P [the system includes H2O(l), O2(g), and O2(aq); you increase the P(O2)]
b. N2(g), increase V [increase the volume of the “system” without adding more moles]
c. He(g), decrease T
d. NaI(s), increase P
12. When you heat water on a stove, bubbles form on the bottom of the pot before the water boils. What are
these bubbles?
13. A student bought a goldfish in a pet shop. Upon returning home, he put the goldfish in a bowl of recently
boiled water that he had cooled quickly. A few minutes later the fish was dead. Why did the student boil the
water? Explain what happened to the fish.
Mole Fraction and Molality
14. Calculate the mole fraction for each component in the following solution: 67.4 g C7H9N in 2.00×102 g
C2H6O. Answ: XC7H9N = 0.127, XC3H6O = 0.873
Unit 9 Student Handout
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15. What is the mole fraction of naphthalene, C10H8, in 0.250 m solution of C10H8 in C6H6? Answ: XC10H8 =
0.0191
16. What is the mole fraction of
a. urea, CO(NH2)2, in a solution prepared by dissolving 25.2 g CO(NH2)2 in 125 mL of water at 20°C (d =
0.998 g/mL) Answ: Xurea = 0.0571
b. ethanol in an aqueous solution that is 22.4% ethanol by mass? Answ: Xethanol = 0.101
17. Calculate the molality of each of these solutions:
a. 14.3 g sucrose (C12H22O11) in 676 g of water Answ: m = 0.0618
b. 7.20 moles of ethylene glycol (C2H6O2) in 3546 g of water. Answ: m = 2.03
18. Calculate the molality of each of the following solutions:
a. 62.1 g urea, (NH2)2CO, in 849 g of water Answ: m = 1.22
b. 1.52 g of sulfur, S8, in 356 g of naphthalene. Answ: m= 0.0166
Unit 9 Student Handout
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19. A 10.7 m solution of NaOH has a density of 1.33 g/cm3 at 20°C. Calculate the following:
a. The mole fraction of NaOH Answ: XNaOH = 0.162
b. The weight percent of NaOH Answ: = 30.0%
c. The molarity of the solution Answ: 9.91 M
20. Determine the grams of solute required to prepare a 3.00 m solution of KOH containing 1500. g of water.
Answ: 252 g KOH
Colligative Properties
Use the symbol > or < to complete each of the following comparisons between vapor pressures, boiling points,
and freezing points of (1) a pure solvent, and (2) a solution of a nonvolatile solute in the solvent at the same
temperature.
1. vapor pressure: pure solvent ______ solution
2. boiling point: pure solvent ______ solution
3. freezing point: pure solvent _______ solution
4. Describe and explain the effect of adding a nonvolatile, non-ionizing solute on the vapor pressure of a pure
solvent.
5. A solution is prepared by dissolving 396 g of sucrose (C12H22O11, 342 g/mol) in 624 g of water. What is the
vapor pressure of this solution at 30°C? (The vapor pressure of water is 31.8 mmHg at 30°C.) Answ: 30.7
mm Hg
Unit 9 Student Handout
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6. Why is the addition of antifreeze, C2H6O2, to the water in a car radiator just as important in hot climates as in
cold climates?
7. Ideally, a 1 m solution of zinc chloride (ZnCl2) in water would boil at 101.545°C, whereas a 1 m solution of
glucose (C6H12O6) would boil at 100.515°C.
a. What is the boiling point elevation of an ideal solution of zinc chloride in water?
b. Compare the boiling point elevation of a zinc chloride solution with the boiling point elevation of a
glucose solution.
c. How does the difference in boiling point elevation provide evidence that zinc chloride dissociates in
water?
Kb water = 0.515°C/m, and Kf water = 1.853°C/m.
8. Calculate the freezing point of a solution that contains 92.0 g of ethanol, dissolved in 500.0 g of H2O. Answ:
-7.41 °C
9. Calculate the boiling point of the solution described in the previous problem. Answ: 102.06 °C
10. What is the boiling point elevation when 67.7 g of urea (NH2CONH2) is dissolved in 833 g of chloroform?
(Kb for chloroform is 3.85°C/m.) Answ: +5.24 °C
Unit 9 Student Handout
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3
11. How many grams of ethylene glycol (antifreeze, C2H6O2) must be dissolved in 2.100 x 10 g of water to
3
lower the freezing point by 20.0°C? Answ: 1.41x10 g
12. If 152 g of sodium sulfate, Na2SO4, are dissolved in 875 g of H2O, what will be the boiling point of the
resulting solution? (Assume 100% ionization.) Answ: 101.88 °C
13. Compute the freezing point of a solution in which 55.6 g of Al(NO3)3 are dissolved in 1250 g of H2O.
(Assume 100% ionization.) Answ: -1.55 °C
14. Calculate the molar mass of a nonionic solute when 7.60 g are dissolved in 475 g of H2O. The resulting
solution has a freezing point of -0.753°C. Answ: 39.4 g/mol
15. Arrange the following aqueous solutions in order of decreasing freezing point, and explain your reasoning:
0.50 m HCl, 0.50 m glucose, 0.50 m acetic acid.
Unit 9 Student Handout
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16. Arrange the following aqueous solutions in order of decreasing freezing point: 0.10 m Na3PO4, 0.35 m
NaCl, 0.20 m MgCl2, 0.15 m C6H12O6, 0.15 m HC2H3O2.
Unit 9 Student Handout
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