Chapter 10 & 17
The Condensed States of Matter
Chapters 10 & 17: The Condensed States of Matter
Notes – Part 1: The Nature of Solids and Liquids
Objectives: Describe the nature of a liquid in terms of
the attractive forces between the representative particles.
Differentiate between evaporation and
boiling of a liquid using the kinetic molecular theory.
Describe how the degree of organization
of particles distinguishes solids form liquids and gases.
Distinguish between crystal lattice
and unit cell.
Explain how allotropes of an element
differ.
Identify, define, and explain: vaporization, evaporation, vapor pressure,
boiling point, normal boiling point, melting point, crystal, unit cell, allotropes,
molecular solids, metallic solids, ionic solids, network (covalent) solids,
and amorphous solids.
Differentiate between the various types of solids using their intermolecular
forces and other properties.
Text Reference: Section 10.2 – pages 274-279 and Section
10.3 – pages 280-283
Initial Question: What makes a liquid different from a
solid? How about the KE of the particles of solids and liquids?
Condensed States:
What is VAPORIZATION?
What is EVAPORATION?
What is the difference between vaporization and evaporation?
Evaporation is a cooling process.
What does this statement mean and why do things evaporate more quickly when
heated?
What do you know about the kinetic energy of the molecules of a sample of
a solid or a liquid?
Do all the molecules have exactly the same kinetic energy?
Equilibrium:
Example of reversible conditions. . .
Vapor Pressure:
Remember temperature is proportional to kinetic energy. T ∝ KE
Question: How do you increase the vapor pressure of a liquid?
Why?
Question: How do you increase the vapor pressure of a solid?
Why?
Question: What is the difference between the vapor (from vapor pressure)
and a gas?
Question: There are three different molecular forms of carbon.
What are these three forms?
Allotrope:
Crystal Lattice:
Unit Solid:
There are various types of solids: amorphous, molecular, metallic,
ionic, and network (covalent).
What is a key difference between these types of solids?
Type (& define) Melting Point Malleability
& Ductility Conductivity Examples
Diagram
Amorphous
Molecular
Metallic
Ionic
Network (Covalent)
Chapters 10 & 17: The Condensed States of Matter
Assignment – Part 1: The Nature of Solids and Liquids
1. Differentiate between boiling and evaporation of a liquid?
2. Why does evaporation lower the temperature of the liquid?
3. How do solids differ from liquids with regard to particle
organization and energy?
4. How does the crystal lattice of a solid differ from
its unit cell?
5. How do the allotropes of an element differ?
6. Explain vapor pressure with regard to dynamic equilibrium.
7. Explain why increasing the temperature of a liquid increases
its rate of evaporation.
8. Would you expect an equilibrium vapor pressure to be
reached above a liquid in an open container? Why?
9. At the top of Mt. Everest, water boils at only 69oC.
Use Figure 10.11 in your text to estimate the atmospheric pressure at the
top of this mountain.
10. Why is the equilibrium that exists between a liquid
and its vapor in a closed container called a dynamic equilibrium?
11. If the volume of the container in which there is a
liquid-vapor equilibrium is changed, the vapor pressure is not affected.
Why?
12. Friend A says that ionic compounds do not conduct electricity.
Friend B says that ionic compounds do conduct electricity. How is it
that each statement has a bit of truth in it?
Chapters 10 & 17: The Condensed States of Matter
Notes – Part 2: Crystalline Lattice and Units Cells & Metallic Bonding
Objectives: Distinguish between a crystal lattice and a
unit cell.
Identify, define, and explain: cubic cell, tetragonal cell, orthorhombic
cell, monoclinic cell. triclinic cell, hexagonal cell, rhombohedral cell,
simple cubic, body-centered cubic, and face-centered cubic.
Describe the arrangement of atoms in some common metallic crystal structures.
Text Reference: Section 10.3 (Part ) – pages 280-283 and
Section 15.3 (Part) – pages 427-429
Read sections 10.3 (pages 280-283) and 15.3 (pages 427-429) in your text.
Answer the questions and complete the Guided Practice Sheets that I have
given you.
Then complete the following assignment.
Chapters 10 & 17: The Condensed States of Matter
Assignment – Part 2: Crystalline Lattice and Units Cells & Metallic Bonding
1. Name at least one physical property that would permit
you to distinguish between a molecular solid and an ionic solid.
2. Describe what happens when a solid is heated to its
melting point.
3. Molecular solids usually have lower boiling points than
covalent network solids. Why?
4. Use Figure 10.14 in your textbook to identify the crystal
structures described by the following characteristics.
a. three unequal axes mutually perpendicular
b. three equal axes making equal angles
with each other
c. two equal axes and one unequal axis
mutually perpendicular
d. three unequal axes intersecting obliquely
e. three axes equal and mutually perpendicular
5. Use metallic bonding theory to explain the physical
properties of metals.
6. Describe the arrangement of atoms in metallic crystal
structures.
7. Why aren’t the properties of all steel samples identical?
Chapters 10 & 17: The Condensed States of Matter
Notes – Part 3: Changes of State and Phase Change Diagrams
Objectives: Interpret the phase diagram of water at any
given temperature and pressure.
Describe the behavior of solids that change directly to the gaseous state
and re-condense to solids without passing through the liquid state.
Identify, define, and describe: phase diagram, triple point, sublimation,
deposition, boiling, evaporation, melting, and phase change.
Text Reference: Section 10.4 – pages 284-286
Question: How are solid, liquid, and vapor states of water related?
Recall:
Equilibrium: the condition, in any reversible process, where the forward
and reverse processes occur at the same rate
Vapor Pressure: the pressure of the vapor over a solid or liquid when
the two states are in equilibrium
Now, you know that when a substance boils, it changes from a liquid to a
gas, but what exactly is boiling point?
Boiling Point:
Graph and Questions: The vapor pressure of a gas above its own (solid
or) liquid depends on temperature. The boiling point or temperature
at which bubbles of vapor form within a liquid, depends on both vapor pressure
and atmospheric pressure. The following table shows the vapor pressure
of a certain liquid at various temperatures. Graph the data on the
grid provided.
Temperature
(oC) Pressure
(kPa)
0 1
10 2
20 4
30 7
40 11
50 16
60 22
70 29
80 37
90 46
100 56
110 68
120 80
130 90
140 107
150 122
1. What effect does increasing the temperature have on
the vapor pressure? Explain in terms of energy and forces.
2. If the atmospheric pressure during this experiment is
96 kPa, what would be the boiling point of the graphed substances?
Mark this point on the graph.
3. What would be the boiling point, if the atmospheric
pressure were to rise? WHY???
4. What would happen to the boiling point if the liquid
were tested at a higher altitude? WHY???
5. How would the time it takes for an egg to fully hard
boil be affected by a higher altitude?
6. Why do we say that water boils at 100oC? What
is 100oC in terms of water?
Normal boiling point:
Phase diagram:
Here is the phase diagram for water.
Image Source: http://wps.prenhall.com/wps/media/objects/602/616516/Chapter_10.html
- Each region represents a phase of water.
- A line that separates two regions shows the conditions
at which those two phases are in equilibrium.
- The curved line between liquid and vapor also shows how
water’s vapor pressure varies with temperature.
Triple Point:
For water, the triple point is:
0.0098oC and 0.0060 atm (0.61 kPa).
Notice, the normal boiling and melting points indicated on the phase diagram.
Question: What happens if boiling and melting are carried out at temperatures
less than 1 atm?
Question: What happens to the boiling point, if the pressure is higher
than 1 atm?
Question: At what temperature and pressure are the liquid and solid
phases of water in dynamic equilibrium?
Volatile:
Volatile substances have high vapor pressure because they
turn easily from a liquid to a gas.
Rubbing alcohol is an example of a highly volatile substance.
What can you say about the intermolecular forces of volatile substance?
Question: Which is more volatile: olive oil or rubbing alcohol?
Question: Which is more volatile: acetone (nail polish remover) or
motor oil?
Chapters 10 & 17: The Condensed States of Matter
Assignment – Part 3: Changes of State and Phase Change Diagrams
1. Explain/Define triple point.
2. What is a practical use of sublimation?
3. When you remove the lid from a food container that has
been left in the freezer for several months, you discover a large collection
of ice crystals on the underside of the lid. Explain why this has happened.
4. The temperature of the gas in an aerosol container is
0oC (273K). to what temperature must the gas be raised to increase
the average kinetic energy of the gas molecules by a factor of three?
5. Both carbon tetrachloride and mercury are liquids at
room temperature and both have vapors that are highly poisonous to humans.
If carbon tetrachloride is spilled in a classroom, the room is aired out
overnight and the threat of harm is eliminated. If mercury is spilled
in a classroom, it is necessary to pick up the liquid droplets with a vacuum
device in order to eliminate the threat of harm. Explain why the precautions
for mercury and carbon tetrachloride are so different since they are both
liquids.
The following table shows the vapor pressure of two liquids at various temperatures.
Graph this information on the grid provided.
Temperature Pressure (kPa) Pressure (kPa)
(oC) Subst. A Subst. B
0 1 1
10 1 2
20 3 4
30 5 7
40 8 11
50 14 16
60 19 22
70 26 19
80 37 37
90 43 46
100 49 56
110 60 68
120 67 80
130 78 93
140 88 107
150 101 122
6. Compare the relative strength of the intermolecular
forces in these two liquids.
7. If the atmospheric pressure is 96 kPa, what is the boiling
point of each substance on the graph?
Substance A ____________________ Substance
B ______________________
Mark and label these points on your graph.
Chapters 10 & 17: The Condensed States of Matter
Notes – Part 4: Heating Curves and Phase Changes
Objectives: Classify, by type, the heat changes that occur
during melting, freezing, boiling, and condensation.
Identify, define, and explain: heat of fusion, heat of vaporization, heat
capacity, heating curve, cooling curve, heat, calorie, and joule.
Graph and interpret a heating curve or cooling curve.
Text Reference: Section 11.3 (Part) – pages 307-313
Question: On a molecular level, what is the difference between
liquid water and water vapor?
Question: Is evaporation endothermic or exothermic?
Question: Because of the attraction between molecules in a liquid,
what must happen for a liquid to change to a gas?
Heat supplied at a constant rate to a substance will be absorbed by that
substance. Temperature menasurements reveal energy changes and changes
of state of the substance. Let’s examine temperature data recorded
when a sample of ice was heated from a temperature of -10oC to a temperature
of 110oC. Graph this information ON THE GRID PROVIDED.
Time
(min) Temperature
(oC)
0 -10
1 -5
2 0
3 0
4 0
5 10
6 20
7 30
8 40
9 50
10 60
11 70
12 80
13 90
14 100
15 100
16 100
17 100
18 100
19 100
20 100
21 105
22 110
Note, there are 5 segments in this graph. It shows the heating of solid
H2O (ice) to gaseous H2O (water vapor). What is happening in each segment?
Segment 1:
Segment 2:
Segment 3:
Segment 4:
Segment 5:
The above graph gives a qualitative picture of the heating of water from
solid to gaseous states. But, in chemistry we need to be quantitative.
Let’s define some key terms.
calorie:
Calorie:
Joule:
Conversion: 1 calorie = 4.184 J
Specific Heat:
Specific Heat Capacity = Cp:
The amount of heat needed to raise the temperature of 1 g of a substance
by 1oC varies from substance to substance. It also varies with the
state of the substance. The specific heat for a substance is different
for its solid, liquid, and gaseous states. Also, when the specific
heat is given, the temperature must be stated, because the specific heat
varies from one temperature to another.
Examine the specific heat for water.
water(s) = 2.060 J/goC at 0oC
water(l) = 4.22 J/goC at 0oC
water(l) = 4.18 J/goC at 25oC
water(g) = 2.070 J/goC at 100oC
water(l) = 4.21 J/goC at 100oC
A substance with a high specific heat can be used as a “heat sponge” and
may be used to absorb and store extra heat.
If the temperature of a substance is increased, the substance has absorbed
energy. If the temperature of a substance is decreased, the substance
has released energy.
The specific heat of water, as listed above, is 4.18 J/goC at 25oC.
This specific heat uses joules as the unit of energy. Another unit
may also be used: calories. Since a calorie is also a unit of energy,
the specific heat of water may also be designated as 1.00 cal/goC at 25oC.
Question: What does a sloped line in a “heating curve”
indicate?
What does a plateau in a “heating curve”
indicate?
Question: Which segment is longer, segment 2 or segment
4? Why do you think this segment is longer?
Question: What would the graph look like from 10 – 21 minutes
if the data was collected in the mountains?
Question: The graph above is a heating curve. What
would the graph of a substance being cooled look like?
Question: Why are there plateaus on the heating curve?
If heat is still added, why is there no increase in temperature?
Chapters 10 & 17: The Condensed States of Matter
Assignment – Part 4: Heating Curves and Phase Changes
1. What is the relationship between kinetic energy and
temperature?
2. A beaker contains water and ice cubes, all at 0oC.
An ice cube with a temperature of –5oC is added. After equilibrium
is re-established, the temperature is still 0oC.
a. What happened to the temperature
of the ice cube that was added?
b. What happened to some of the water
after the ice cube is added?
c. A small amount of warm water is added to the beaker,
but the equilibrium is still 0oC. What happened to the warm water?
d. What happened to some of the ice
after the warm water is added?
e. What happened to some of the original
0oC water after the warm water is added?
3. You are camping in Death Valley. What happens
when you go to boil water for pasta? How long does it take to get to
the boiling point? What happens to the length of time it takes for
the pasta to cook?
4. How would the time it takes for an egg to fully hard
boil be affected by a higher altitude?
5. In an experiment, a round bottom flask is evaporated
to near zero pressure. Some liquid alcohol is injected into the flask.
After the pressure gauge reads a stabilized pressure, there is still some
alcohol left in the flask. The attached pressure gauge indicates a
pressure of 0.733 atm.
a. What is the vapor pressure of the alcohol?
b. Since the flask has no initial pressure, from where
did the resulting pressure come?
c. What would happen to the pressure reading if the temperature
of the alcohol were increased?
d. What would happen to the amount of liquid alcohol if
the temperature of the alcohol were increased?
6. Why is a burn from steam potentially far more serious
than a burn from very hot water?
7. Why does an ice cube melt at room temperature?
8. Equal masses of two substances absorb the same amount
of heat. The temperature of substance A increases twice as much as
the temperature of substance B. Which substance has the higher specific
heat? Explain.
Chapters 10 & 17: The Condensed States of Matter
Notes – Part 5: Heat Calculations
Objectives: Calculate heat changes that occur during heating,
melting, freezing, boiling, condensing, and cooling.
Identify, define, and explain: specific
heat capacity, heat of vaporization and heat of fusion.
Text Reference: Section 11.3 – pages 307-313
Recall, specific heat capacity (Cp) is the amount of heat required to raise
the temperature of 1 g of a substance by 1oC.
On what part of a heating curve would specific heat capacity be significant?
Heat quantities depend on the substance’s specific heat, the quantity being
heated, and the temperature change.
The heat change, resulting in a temperature change, may be calculated for
the following: a solid being heated to its melting point, a liquid being
heated to its boiling point, a gas being superheated, and any of the reverse
processes for the substance being cooled.
The formula for calculating heat changes resulting in a temperature change:
__________ = ______________________________ with a unit of _________________________
__________ = ______________________________ with a unit of _________________________
__________ = ______________________________ with a unit of _________________________
__________ = ______________________________ with a unit of _________________________
Example 1: How much heat is absorbed by 126 g of
water if its initial temperature is 47.9oC and its final is 75.6oC?
Question: Can this formula be used for a sample of
water with a temperature change from 56.9oC to 123.9oC? Explain.
The specific heat of a substance describes the energy that must be added
to increase the temperature of a substance or the amount of energy that must
be removed to lower its temperature. However, when you change the temperature
of a substance, it may undergo a change of state, thereby requiring the use
of a different specific heat. Also, there are tremendous amounts of
energy involved in changes of state.
Molar heat of fusion:
Molar heat of vaporization:
So if there is no temperature change and the change occurring is a change
of state, you need to calculate the energy using:
Q = m Cfus or Q = m Cvap
Example 2: How much energy, in joules, is required
to boil 75.8 g of water at 100oC?
When you solve a heat problem where the substance passes through a phase
change or multiple phase changes, you will need to complete multiple steps
to arrive at the final answer.
Use the following values for H2O: H2O
– solid Cp = 2.060 J/goC
Cfus = 334 J/g
H2O – liquid
Cp = 4.180 J/goC
H2O – gas Cp = 2.070 J/goC
Cvap = 2260 J/g
Example 3: You have a sample of ice at –47oC and the sample
has a mass of 357.9 g.
You need to get the temperature of your sample to +118.6oC.
What is the total amount of heat required for this change? (Show each
part separately!!!)
In this example, you are starting at a temperature below the freezing point
and you need to get to a temperature above the freezing point. Each
segment on the change of state graph needs to be completed separately.
You will calculate the five steps: below freezing to freezing, from solid
to liquid, from freezing to boiling, from liquid to gas, from boiling to
above boiling. After each of the steps has been calculated, you will
add together each portion to arrive at your final answer.
Below freezing point to freezing point:
At freezing point – from a solid to a liquid:
From freezing point to boiling point:
At boiling point – from a liquid to a gas:
From boiling point to above boiling point:
Total heat involved in the change:
Careful to keep ΔT = Tfinal - Tinitial
What does it mean if you get a negative value for ΔT?
Why is it important to keep the negative sign for ΔT?
Chapters 10 & 17: The Condensed States of Matter
Notes – Part 5: Heat Calculations
Solve the following problems. Be sure to show everything you need to
show to receive complete credit.
1. How much energy, in joules, is required to freeze 250.0
g of water to ice at 0.0oC?
2. A 325.0 g sample of water loses 765.4 J of heat.
What temperature change did it undergo?
3. A sample of water is cooled from 76.5oC to 43.2oC and
543.2 J of heat is released. What is the mass of the sample?
4. A sample of an unknown substance has a mass of 54.8
g. Heat is added and the temperature increases from 35.75oC to 58.25oC.
If 3256 J of heat are required for this temperature change to occur, what
is the specific heat capacity of this substance.
5. The same amount of heat is added to a ton of iron and
a small iron nail, which were originally at the same temperature.
a. Which would reach a higher temperature?
b. Why is there a difference?
Explain in terms of concentration or intensity of heat in matter.
6. A cup of boiling water has less heat than a large iceberg.
Explain why this is true.
7. The Cp of ammonia gas is 2.190 J/goC. If you have
a sample of 135.7 g of ammonia gas that absorbed 3527 J of heat, what temperature
change has the ammonia gas has undergone?
8. Use the following values fro substance X:
X – solid Cp = 1.356 J/goC
Cfus = 264 J/g
X – liquid Cp = 3.498 J/goC
X – gas Cp = 1.549
J/goC Cvap = 984 J/g
X – boiling point = 82.467oC
X – freezing point = -12.360oC
You have a sample of X with a mass of 168.26 g at an initial temperature
of –35.68oC.
You want to get the sample to a final temperature of 91.03oC.
What is the total heat, in joules, required for this temperature change to
occur? (Show each part separately.)
Chapters 10 & 17: The Condensed States of Matter
Notes – Part 6: Water – An Unusual Substance
Objectives: Describe the hydrogen bonding that occurs in
water and represent in graphically.
Explain the high surface tension and
low vapor pressure of water in terms of hydrogen bonding.
Identify, define, and explain: surface
tension and surfactant.
Account for the high heat of vaporization
and the high boiling point of water in terms of hydrogen bonding.
Explain why ice floats in water.
Text Reference: Section 17.1 – pages 475-478 and Section
17.2 – pages 479-481.
Read sections 17.1 (pages 475-479) and 17.2 (pages 479-481) in your text.
Answer the questions and complete the Guided Practice Sheets that I have
given you.
Then complete the following assignment.
Chapters 10 & 17: The Condensed States of Matter
Assignment – Part 6: Water
1. Describe hydrogen bonding between molecules in water?
2. How is hydrogen bonding responsible for the high surface
tension of water and the low vapor pressure of water?
3. What is a surfactant?
4. What happens to a bottle of soda that you left in the
freezer because you forgot about it?
5. Examine Table 17.2 on page 480 in your text book.
The table shows a trend for the density of liquid water to increase with
decreasing temperature. Why does this trend not continue below 4oC?
6. Why does water have a relatively high boiling point
and heat of vaporization?
7. What is the difference between the structure of liquid
water and the structure of ice? How does this explain why ice floats
in water?