Chapter 13 States of Matter
FOOD FOR THOUGHT…

What is the relationship between solids,
liquids, and gases? How are they the same?
How are they different?

Make a table or list answering the above question.
You will need your answers tomorrow.
13.1 - Gases
1.
2.
Flemish physician Jan
Baptista Van Helmont
used the Greek word
chaos, meaning
without order, to
describe the products
or reactions that had
no fixed shape or
volume.
From chaos came the
term gas.
13.1 - Gases
3.
4.
Around 1860, Ludwig
Boltzmann and James
Maxwell each proposed
a model to explain the
properties of gases.
The model is known as
the kinetic-molecular
theory.
Kinetic-molecular theory
describes the behavior
of gases in terms of
particles in motion.
13.1 - Gases
Kinetic – molecular theory: (pages 386, 419-420)
5.
Gas particles do not attract or repel each other.
1.

Why?
Gas particles are much smaller than the distance
between them. The theory assumes the
particles have virtually no volume.
2.

What process can be applied to gases because of the very
small particle size?
Gas particles are in constant, random motion.
Because of this, gas particles spread out and
mix.
3.

How do gas particles move? What may cause their motion
to change?
13.1 - Gases
Kinetic – molecular theory: (pages 386, 419-420)
5.
No kinetic energy is lost when gas particles collide
with each other or with the walls of their container.
4.



Why?
How can you describe elastic collisions?
What can be said about the total kinetic energy of a
container of gas?
All gases have the same average kinetic energy at a
given temperature.
5.




What happens to total energy of a gas as temperature
increases?
What happens to total energy of a gas as temperature
decreases?
What two factors determine kinetic energy?
What is temperature? What do all gases have in common?
Food for thought…
Will the kinetic-molecular theory
work for liquids and solids?
Explain your reasoning.
13.1 - Gases
6.
1.
How does kinetic-molecular theory explain the behavior
of gases?
Density

The idea that there is a lot space between particles explains
the low density (mass / volume) that gases have. There are
fewer molecules of gas in a given amount of volume.
Compression and Expansion
2.




Gas molecules can be compressed (pressed into a smaller
space using a piston).
Gas molecules can expand as pressure from the piston is
released.
What happens to the density as a gas is compressed?
What happens to the density as a gas expands?
13.1 - Gases
6.
3.
How does kinetic-molecular theory explain the behavior
of gases?
Diffusion





The idea that there are no significant forces of attraction
between gas particles supports the fact that gases can flow
easily past each other.
When gases mix, what will eventually happen to their
concentration?
What is diffusion?
In which direction does diffusion occur?
What does the rate of diffusion depend upon?
Effusion
4.


What is effusion?
How can effusion be compared to diffusion?
13.1 - Gases
6.
4.
How does kinetic-molecular theory explain the behavior
of gases?
Effusion


Graham’s law of effusion – the rate of effusion for a gas is
inversely proportional to the square root of its molar mass
Graham’s law applies to diffusion rates also.
Rate of effusion
Rate A
Rate B
1
1
SQRT(molar mass)
= SQRT(molar mass B / molar mass A)
Food For thought…
1. What state of matter is He?
2. What is happening to the He atoms inside
a balloon to keep the balloon inflated?
Use the kinetic-molecular theory to support
your answers.
13.1 - Gases
TIME TO WORK!
Effusion / diffusion problems (1-2), page 388
Dalton’s Law problems (4-6), page 392
Problem-solving lab, page 390
Calculate the ratio of effusion rates for
pairs of the noble gases.
13.2 – Forces of Attraction
7.
8.
Intramolecular Forces

The forces that hold particles together in
ionic, covalent, and metallic bonds are called
intramolecular forces. These are forces
occurring within the chemical compound.
Intermolecular Forces

Intermolecular forces happen between or
among like molecules of a substance.

There are 3 intermolecular forces we will
discuss: dispersion forces, dipole-dipole
forces, and hydrogen bonds.
13.2 – Forces of Attraction
9.
Intramolecular Forces

All intermolecular forces are weaker than
intramolecular bonding forces.
10. Relative Strength of Molecular Forces
Covalent network > ionic bonds > metallic bonds >
Hydrogen bonds > dipole-dipole forces > London dispersion
forces
Food for thought…
Many words can be understood by looking at
their parts.
 Match the following word parts to their
meanings.
Heat
-ion
thermThe result of an action
-ic
-ize
Related to
endoTo become
Inside
exo
Outside
13.2 – Forces of Attraction
11.
a.
b.
c.
d.
e.
f.
London dispersion forces –
are weak forces that result from temporary
shifts in the density of electrons in electron
clouds (draw sketch page 394)
are named after the German-American physicist
who first described them, Fritz London
are weak forces because they are based on
temporary dipoles
are the dominant force of attraction between
identical nonpolar molecules
dispersion force strength: I > Br > Cl > F
explains why F and Cl are gases, Br is liquid,
and I is solid at room temperature
13.2 – Forces of Attraction
12.
a.
b.
c.
d.
e.
f.
Dipole-dipole forces –
are forces of attraction between oppositely charged
regions of polar molecules.
are weak forces of attraction that result from permanent
dipoles within polar molecules
Some regions of a polar molecule are always partially
negative and other regions are always partially positive.
Neighboring polar molecules orient themselves so that
oppositely charged regions line up. (sketch diagram page
394)
The degree of polarity in a molecule depends on the
relative electronegativity values of the elements in the
molecule.
Dipole forces are stronger than dispersion forces as long
as the molecules being compared have about the same
mass.
Bell ringer…



Which are stronger, intermolecular or
intramolecular forces?
Which specific force is strongest?
Which is weakest?
13.2 – Forces of Attraction
13.
a.
b.
c.
d.
e.
Hydrogen bonds –
are a special type of dipole attractive force between
highly polar molecules.
are dipole-dipole attractions that occur between
molecules containing a hydrogen atom bonded to a small,
highly electronegative atom with at least one lone
electron pair. (sketch diagram page 395)
Hydrogen must be bonded to either a F, O, or N atom
because these atoms are electronegative enough to
cause a large partial positive charge on the hydrogen, yet
are small enough that their lone pairs of electrons can
come close to hydrogen atoms.
In hydrogen bonds, hydrogen atoms on one molecule of a
substance are attracted to the negative end of another
molecule of the substance.
Hydrogen bonds explain why water is liquid at room temp.
13.3 – Solids and Liquids
14.
Although kinetic-molecular theory was
developed to explain the behavior of gases, the
model can be applied to liquids and solids.
However, you must consider the forces of
attraction between particles as well as their
energy of motion.
In liquids and solids, particles are closer together
Why consider the force of attraction between
because of stronger forces of attraction between them. In
particles
of liquids
and solids?
liquids,
the forces
of attraction
limit their range of motion
so that particles are closely packed in a fixed volume.
In solids, the forces are so strong that motion is limited to
vibrations around a fixed location. Therefore, solids have
more order in their particles.
Food for thought…

Put the following attractive forces in the
proper category: intermolecular or
intramolecular
dispersion forces
dipole-dipole
metallic bonds
hydrogen bonds

covalent bonds
ionic bonds
Put the attractive forces in order from weakest to
strongest.
13.3 – Solids and Liquids
Liquids - define the following terms related
to properties of liquids
15.
1.
2.
3.
4.
5.
6.
7.
8.
Density
Compression
Fluidity
Viscosity
Temperature (relationship to viscosity)
Surface tension (define and give an example)
Surfactants
Capillary action (define and give an example)
13.3 – Solids and Liquids
Solids - define the following terms related
to properties of solids
16.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
Density
Compressibility
Fluidity
Crystalline solid
Unit cell
Molecular solids
Covalent network solids
Ionic solids
Metallic solids
Amorphous solids
13.3 – Solids and Liquids

Sketch and label 3 basic solid crystal lattice
structures (page 400).
13.4 – Phase Changes
17.
6 possible transitions between phases
GAS
MELTING
SOLID
FREEZING
LIQUID
13.4 – ENDOTHERMIC CHANGES
18.
Phase changes requiring energy input
a.
Melting





What is melting?
What is heat?
Describe what happens when ice melts.
Why is the energy required to melt salt
much greater than the energy to melt ice?
What is melting point?
Endothermic Phase Changes
Phase changes requiring energy input
Melting
Melting point
is when
of aenergy
solid isadded
the temperature
to a solid isat
great
which
enough
forcesto
break the
holding
itsforces
structure
holding
together
the are
molecules
brokentogether.
and it becomes a
The solid
liquid
.
phase changes to a liquid.
Heat is the transfer of energy from a higher temperature to a
lower temperature.
When ice melts, molecules on the surface absorb enough
Energy to break the hydrogen bonds. The molecules move
apart and become a liquid.
The ionic bonds in salt (sodium chloride) are much stronger
than the hydrogen bonds in ice.
13.4 – ENDOTHERMIC CHANGES
18.
Phase changes requiring energy input
b.
Vaporization






When can the temperature of a melting
substance start to rise?
How much energy does a substance need
to vaporize?
What is a vapor?
What is vaporization?
What is evaporation?
How long does it take for evaporation to
occur?
Endothermic Phase Changes
Phase changes requiring energy input
Vaporization
Evaporation
The
temperature
is vaporization
of a melting
that
substance
occurs only
will at
start
thetosurface
rise after
all athe
of
liquid
solid
. substance has melted.
To
a substance depends
needs enough
energy
to overcome
Thevaporize,
time for evaporation
upon the
amount
of
the
forces
of attraction
holding
molecules
together in the
liquid
and the
amount of
energythe
available
.
liquid.
A vapor is a substance that is ordinarily a liquid at room
temperature.
Vaporization is the process by which a liquid changes to a gas
or vapor.
13.4 – ENDOTHERMIC CHANGES
18.
Phase changes requiring energy input
b.
Vaporization




What is vapor pressure?
How does a rise in temperature affect vapor
pressure?
What is the boiling point?
As temperature increases, what happens to
kinetic energy of molecules?
Endothermic Phase Changes
Phase changes requiring energy input
Vaporization
Vapor pressure is the pressure exerted by a vapor over a
liquid.
Vapor pressure increases with increasing temperature.
The boiling point is the temperature at which the vapor
pressure of a liquid equals the outside or atmospheric
pressure.
Kinetic energy of molecules increases with temperature.
Endothermic Phase Changes
Phase changes requiring energy input
Sublimation


What is sublimation?
Name some substances that sublime.
Sublimation is the process by which a solid changes directly
to a gas without becoming a liquid.
Solid iodine, solid carbon dioxide (dry ice), moth balls,
air fresheners, ice cubes left in freezer.
13.4 – Exothermic Changes
Phase changes that release energy
19.
Condensation
a.





Name some examples of condensation.
What is condensation?
Describe what happens to molecules when
they condense.
What happens when hydrogen bonds form
in liquid water?
Name some circumstances that will cause
condensation.
Exothermic Phase Changes
Phase changes that release energy
Condensation
Examples:
dew,
frost,
water on
of a glass, water on
Condensation
can
be caused
byoutside
vapor molecules
a
window pane,
air conditioning
condensate
contacting
a coldclouds,
surfacefog,
(drink
glass, dew) or
cold air (fog,
clouds).
Condensation is the process by which a gas or a vapor
becomes a liquid. It is the reverse of vaporization.
When vapor molecules condense, they lose energy, slow
down, and form bonds with each other when they collide.
The bonded molecules are more dense and become a
liquid.
When hydrogen bonds are formed in water, energy is released.
Exothermic Phase Changes
Phase changes that release energy
Deposition


What is deposition?
Give some examples of deposition.
Deposition is the process by which a substances changes
from a gas or vapor to a solid without first becoming a liquid.
It is the reverse of sublimation.
Example: snow
13.4 – Exothermic Changes
Phase changes that release energy
19.
Freezing
c.



What is freezing?
Describe what happens when something
freezes.
What is freezing point?
Exothermic Phase Changes
Phase changes that release energy
Freezing
Freezing is the phase change of a liquid to a solid.
During freezing, heat is removed from a substance, the
molecules lose kinetic energy and slow down. When enough
energy has been removed, the molecules become fixed in a
set position.
The freezing point is the temperature at which a liquid is
converted into a crystalline solid. The freezing point and
melting point temperatures are the same for a substance.
15. SUMMARIZING
_______thermic
Solid  _______
________thermic
_______  Solid
_______  Gas
Gas  ______
Liquid  ______
______  Solid
Order of the molecules is
__________________.
Order of the molecules is
__________________.
SUMMARIZING
16.
17.
18.
ENDOTHERMIC
Molecules are_________
gaining energy during an
endothermic phase change.
Kinetic energy – the
What kind of energy? energy of motion
Describe the change in motion of molecules
during melting, vaporization, and sublimation.
The molecules gain energy and
move faster and further apart.
SUMMARIZING
EXOTHERMIC
19.Molecules are__________
losing kinetic energy
during an exothermic phase change.
20.Describe the change in motion of molecules
during freezing, condensation, and deposition.
The molecules lose energy and
slow down getting closer
together.
13.4 – Phase Changes
Phase diagrams
20.
a.
b.
c.
d.
Temperature and pressure are the two
variables that combine to control the phase
of a substance.
A phase diagram is a graph of pressure
versus temperature that shows in which
phase a substance exists under different
conditions of temperature and pressure.
What is the triple point?
All 6 phase changes can occur at the triple
point.
13.4 – Phase Changes
Phase diagrams
20.
e.
What is the critical point?
Problem solving! Add to your tools!
1.
Read the problem statement thoroughly.
2.
Identify the “givens” and write them down.
3.
Identify the “unknowns” and write them down.
4.
Determine the scientific principles you will use.
5.
Devise a strategy for attacking the problem.
Scientific Principles

Name principles you might use to solve various
types of physical science problems.







Law of Conservation of Mass
Rules for Balancing Formulas and Equations
Periodic Table
Solving Mathematical Equations, like F = ma, velocity,
acceleration, momentum
Rules for Balanced and Unbalanced Forces
Gravity – free fall, weight, acceleration due to gravity
Newton’s Laws
Strategies you have used to solve problems…






Using graphs to show data.
Drawing diagrams to show atomic structure.
Using the Periodic Table to determine trends in
atomic radius and ionization energy.
Balancing chemical equations.
Using equations to calculate an unknown value.
Drawing a force diagram and showing the
magnitude and direction of the forces acting on
an object.