The document discusses the kinetic theory description of gases and liquids, explaining that gas particles move rapidly in straight lines and collisions between them result in gas pressure, while liquid particles are able to flow but interact strongly and have less space between them than gas particles. It also covers evaporation as the process where some liquid particles gain enough kinetic energy to overcome attractive forces and enter the gas phase, and how heating increases the rate of evaporation by raising the average kinetic energy of particles.

1. Chapter
13
(Chapter 10 in your book)
“States of Matter”

2. Section 13.1
The Nature of Gases
OBJECTIVES:
 Describe the assumptions of the “kinetic theory” as it applies to gases.
 Interpret gas pressure in terms of kinetic theory.
 Interpret gas pressure in terms of kinetic theory.
 Define the relationship between Kelvin temperature and average kinetic
energy.
 Kinetic refers to ____________________
 The energy an object has because of it’s motion is called ____________________
 The kinetic theory states that the tiny particles in all forms of matter are in
__________________________________________
 Three basic assumptions of the kinetic theory as it applies to gases:
#1. Gas is ______________________________________- usually molecules or atoms
 Small, hard spheres
 Insignificant volume; relatively far apart from each other
 No attraction or repulsion between particles
#2. Particles in a gas move rapidly in ________________________________ motion
 Move in straight paths, changing direction only when colliding with one
another or other objects
 Average speed of O2 in air at 20 oC is an amazing 1700 km/h!
#3. Collisions are ______________________________________- meaning kinetic
energy is transferred without loss from one particle to another- the total kinetic energy
remains constant
 Gas Pressure – defined as ____________________________________________
____________________________________________________________________
 Due to:
a) ___________________________________________
b) ___________________________________________
 No particles present? Then there cannot be any collisions, and thus no
pressure – called a ________________________________
 Atmospheric pressure results from the collisions of air molecules with objects
 Decreases as you climb a mountain because the air layer thins out as
elevation increases
 ________________________________ is the measuring device for atmospheric
pressure, which is dependent upon weather & altitude
2

3. Measuring Pressure
The first device for measuring atmospheric
pressure was developed by Evangelista Torricelli
during the 17th century.
The device was called a “barometer”
• Baro = _________________
• Meter = ________________
 The SI unit of pressure is the _______________________________
 At sea level, atmospheric pressure is about 101.3 kilopascals (kPa)
 Older units of pressure include millimeters of mercury (mm Hg), and
atmospheres (atm) – both of which came from using a mercury barometer
 Mercury Barometer – Fig. 10.2, page 269 – a straight glass tube filled with Hg,
and closed at one end; placed in a dish of Hg, with the open end below the surface
 At sea level, the mercury would rise to 760 mm high at 25 oC- called one
______________________________________ (atm)
 Equal pressures:1 atm = 760 mm Hg = 101.3 kPa
Example Problem
A gas is at a pressure of 1.50 atm. Convert this pressure to a) kilopascals b) millimeters
of mercury
 Most modern barometers do not contain mercury- too dangerous
 These are called aneroid barometers, and contain a sensitive metal
diaphragm that responds to the number of collisions of air molecules
 For gases, it is important to relate measured values to standards
 Standard values are defined as a temperature of 0 oC and a pressure of
101.3 kPa, or 1 atm
 This is called Standard Temperature and Pressure, or ____________
3

4.  What happens when a substance is heated? ______________________________
____________________________________________________________________
_
 Some of the energy is stored within the particles- this is potential energy,
and does not raise the temperature
 Remaining energy speeds up the particles (increases average kinetic
energy)- thus ________________________________________________
 The particles in any collection have a wide range of kinetic energies, from very
low to very high- but most are somewhere in the middle, thus the term
________________________________________________________is used
 The higher the temperature, the wider the range of kinetic energies
 An increase in the average kinetic energy of particles causes the temperature to
rise.
 As it cools, the particles tend to move more slowly, and the average K.E.
declines.
 Is there a point where they slow down enough to stop moving?
 The particles would have no kinetic energy at that point, because they would have
no motion
 ________________________________ ( ____K, or __________ oC) is
the temperature at which the motion of particles theoretically ceases
 This has never been reached, but about 0.5 x 10-9 K has been achieved
 The Kelvin temperature scale reflects a __________________________________
between temperature and average kinetic energy
 Particles of He gas at 200 K have twice the average kinetic energy as
particles of He gas at 100 K
 Solids and liquids differ in their response to temperature
 However, at any given temperature the particles of all substances,
regardless of their physical state, have the same average kinetic energy
 What happens to the temperature of a substance when the average kinetic energy
of its particles decreases?
4

5. Name _______________________________________ Date ______________________
13-1 Section Review
1. According to the assumptions of kinetic theory, how do the particles in a gas
move?
2. Use kinetic theory to explain what causes gas pressure.
3. Express the pressure 545mm Hg in kilopascals.
4. How do you raise the average kinetic energy of the water molecules in a glass of
water?
5. A cylinder of oxygen gas is cooled from 300 K to 150 K. By what factor does the
average kinetic energy of the oxygen molecules in the cylinder decrease?
5

6. Section 13.2
The Nature of Liquids
OBJECTIVES:
 Identify factors that determine physical properties of a liquid.
 Define “evaporation” in terms of kinetic energy.
 Describe the equilibrium between a liquid and its vapor.
 Identify the conditions at which boiling occurs.
 Liquid particles are also in motion.
 Liquid particles are free to slide past one another
 Gases and liquids can both FLOW
 However, liquid particles ______________________________________
to each other, whereas gases are not
 Particles of a liquid spin and vibrate while they move, thus contributing to their
average kinetic energy
 But, most of the particles ________________________ have enough
energy to escape into the gaseous state; they would __________________
___________________________ their intermolecular attractions with other
particles
 The intermolecular attractions also reduce the amount of space between particles
of a liquid
 Thus, liquids are more __________________ than gases
 Increasing pressure on liquid has hardly any effect on it’s volume
 Increasing the pressure also has little effect on the volume of a solid
 For that reason, liquids and solids are known as the
___________________________________________________________
_
 Water in an open vessel or puddle eventually goes into the air
 The conversion of a liquid to a gas or vapor is called _______________________
 When this occurs at the surface of a liquid that is not boiling, the process
is called ________________________________
 Some of the particles break away and enter the gas or vapor state; but only
those with the minimum kinetic energy
 A liquid will also evaporate faster when heated
 Because the added heat increases the average kinetic energy needed to
overcome the attractive forces
 But, evaporation is a __________________________________________
 Cooling occurs because those with the highest energy escape first
6

7.  Particles left behind have _______________________ average kinetic energies;
thus the temperature _________________________
 Similar to removing the fastest runner from a race- the remaining runners
have a lower average speed
 Evaporation helps to keep our skin cooler on a hot day, unless it is very humid on
that day. Why?
 Evaporation of a liquid in a closed container is somewhat different
 When some particles do vaporize, these collide with the walls of the
container producing ___________________________________________
 Eventually, some of the particles will return to the liquid, or _________________
 After a while, the number of particles evaporating will equal the number
condensing- the space above the liquid is now saturated with vapor
 A dynamic equilibrium exists
 Rate of evaporation = rate of condensation
 Note that there will still be particles that evaporate and condense
 But, there will be no NET change
 An __________________________________________________ of a contained
liquid increases the vapor pressure- the particles have an increased kinetic energy,
thus more minimum energy to escape
 The vapor pressure of a liquid can be determined by a device called a
“manometer”- Figure 10.2, p.277
 The vapor pressure of the liquid will push the mercury into the U-tube
 A barometer is a type of manometer
 We now know the rate of evaporation from an open container increases as heat is
added
 The heating allows larger numbers of particles at the liquid’s surface to
overcome the attractive forces
 Heating allows the average kinetic energy of all particles to increase
 The boiling point (bp) is the temperature at which the
______________________
____________________________________________________________________
____________________________________________________________________
__
 Bubbles form throughout the liquid, rise to the surface, and escape into the
air
 Since the boiling point is where the vapor pressure equals external pressure, the
bp changes if the external pressure changes
7

8.  Normal boiling point- defined as the bp of a liquid ________________________
____________________________________________________________________
_
 Normal bp of water = ______________
 However, in Denver = ______________, since Denver is 1600 m above
sea level and average atmospheric pressure is about 85.3 kPa (Recipe
adjustments?)
 In pressure cookers, which reduce cooking time, water boils above 100 oC
due to the increased pressure
 Autoclaves, devices often used in the past to sterilize medical instruments,
operated much in a similar way – higher pressure, thus higher boiling point
 Boiling is a __________________________________ much the same as
evaporation
 Those particles with highest KE escape first
 Turning down the source of external heat drops the liquid’s temperature below the
boiling point
 Supplying more heat allows particles to acquire enough KE to escape- the
_________________________________________________________, the liquid
only boils at a faster rate
 The heat of vaporization of a liquid ____________________________________
____________________________________________________________________
____________________________________________________________________
__
 For water at its normal boiling point of 100 degrees Celcius, the heat of
vaporization is __________________________
 The heat of condensation is __________________________________________
____________________________________________________________________
____________________________________________________________________
____________________________________________________________________
___
 Condensation is exothermic
 For water, the heat of condensation is is 2259 Joules per gram – the same
as the heat of vaporization
8

9. Name _______________________________________ Date ______________________
13-2 Section Review
1. Describe the nature of liquids. Refer to the role of attractive forces in your
answer.
2. Use kinetic theory to explain the differences between the particles in a gas and
those in a liquid
3. Use kinetic theory to explain the differences between evaporation and boiling of a
liquid
4. Explain why the boiling point of a liquid varies with atmospheric pressure
5. Why does evaporation lower the temperature of a liquid?
6. What is the difference between heat of vaporization and the heat of condensation?
9

10. Section 13.3
The Nature of Solids
OBJECTIVES:
 Evaluate how the way particles are organized explains the properties of
solids.
 Identify the factors that determine the shape of a crystal.
 Explain how allotropes of an element are different.
 Particles in a liquid are relatively free to move
 Solid particles are not
 solid particles tend to ___________________________________________,
rather than sliding from place to place
 Most solids have particles packed against one another in a highly organized
pattern
 Tend to be dense and incompressible
 Do not flow, nor take the shape of their container
 Are still able to move, unless they would reach absolute zero
 When a solid is heated, the particles vibrate more rapidly as the kinetic energy
increases
 The organization of particles within the solid breaks down, and eventually
the solid melts
 The melting point (mp) is ____________________________________________
 At the melting point, the disruptive vibrations are strong enough to overcome the
interactions holding them in a fixed position
 Melting point can be reversed by cooling the liquid so it freezes
 Solid ↔ liquid
 Generally, most ionic solids have ______________________________________,
due to the relatively strong forces holding them together
 Sodium chloride (an ionic compound) has a melting point = 801 oC
 Molecular compounds have relatively low melting points
 Hydrogen chloride (a molecular compound) has a mp = -112 oC
 Not all solids melt- wood and cane sugar tend to decompose when heated
 Most solid substances are crystalline in structure
10

11.  When a liquid is cooled, a temperature is eventually reached at which the liquid
begins to freeze. It changes into a solid
 This temperature, which remains constant until all the liquid has solidified at 1
atmosphere pressure, is called the ____________________________________ of
the liquid
 While the liquid is cooling, the average kinetic energy of its particles decreases
until it is low enough for the attractive forces to be able to hold the particles in the
fixed positions characteristic of the solid phase
Heat of solidification
 The amount of heat needed to change ___________________________________
____________________________________________________________________
____________________________________________________________________
__
 Heat of solidification = 334 Joules
Heat of fusion
 The amount of heat needed to change a unit mass of a substance from solid to
liquid at STP
 Heat of fusion of ice = 334 Joules
Heating and cooling curves
us
lsi
Ce 200
in 180
e 160
140
ur
rat
pe 120
m 100
Te 80
60
40
20
0
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16
11

12. Time in minutes
Name _______________________________________ Date ______________________
13-3 Section Review
1. Explain the nature of solids and tell why they differ from liquids. Refer to the
organization of particles in your answer.
2. Would you expect MgCl2 to have a high or low melting point? Explain your
answer.
3. What is the difference between the heat of fusion and the heat of solidification?
4. Describe what happens when a solid is heated to its melting point.
12

13. Section 13.4
Changes of State
OBJECTIVES:
 Identify the conditions necessary for sublimation.
 Describe how equilibrium conditions are represented in a phase diagram.
 Describe how equilibrium conditions are represented in a phase diagram.
 Sublimation- ______________________________________________________
____________________________________________________________________
____________________________________________________________________
__
 Examples: iodine, dry ice (-78 oC); mothballs; solid air fresheners
 Sublimation is useful in situations such as freeze-drying foods- such as by
freezing the freshly brewed coffee, and then removing the water vapor by a
vacuum pump
 Also useful in separating substances - organic chemists use it separate mixtures
and purify materials
 The relationship among the solid, liquid, and vapor states (or phases) of a
substance in a sealed container are best represented in a single graph called a
phase diagram
 Phase diagram- _____________________________________________________
____________________________________________________________________
____________________________________________________________________
__
 The diagram below shows the phase diagram for water
 Each region represents a pure phase
 Line between regions is where the two phases exist in equilibrium
 _________________________________ is where all 3 curves meet, the
conditions where all 3 phases exist in equilibrium
13

14. a)
P Critical
(k Point
re
su
es
Pr
Temperature (oC)
14

15. Name _______________________________________ Date ______________________
13-4 Section Review
1. What general information can you get from a phase diagram for water at various
temperatures and pressures?
2. Describe the process of sublimation. What is a practical use of this process?
3. Explain triple point.
15