Chapter 3: States of
Matter
•Section 3.1 Solids, Liquids, and
Gases
•Section 3.3 Phase Changes
Section 3.1 Solids, Liquids, and
Gases
• Describing the States of Matter
• Key Concept: Materials can be classified
as solids, liquids, or gases based on
whether their shapes and volumes are
definite or variable
• Shape and volume are clues to how the
particles within a material are arranged.
Section 3.1 Solids, Liquids, and
Gases
• Kinetic Theory
• Under ordinary conditions, why are objects
solid, liquid, and gas?
• Kinetic energy-energy an object has due
to its motion (**The faster an object
moves, the more kinetic energy it has.)
• Key Concept: The kinetic theory of matter
says that all particles of matter are in
constant motion.
Section 3.1 Solids, Liquids, and
Gases
• Solids
• Def.-the state of matter in which materials have
a definite shape and a definite volume
• Definite-the shape and the volume of the object
won’t change (does not mean they can never
change)
• **Almost all solids have some type of orderly
arrangement of particles at the atomic level.
Section 3.1 Solids, Liquids, and
Gases
• SOLIDS
Have definite (or fixed) shape and volume
The particles in a solid are held fairly rigidly in place.
Section 3.1 Solids, Liquids, and
Gases
• Explaining the Behavior of Solids
• Key Concept: Solids have a definite volume and
shape because particles in a solid vibrate
around fixed locations.
• **Strong attractions among atoms restrict their
motion and keep each atom in a fixed location
relative to its neighbors.
• Atoms vibrate around their locations but do not
exchange places with neighboring atoms.
Section 3.1 Solids, Liquids, and
Gases
• Liquids
• Def.-the state of matter in which a material
has a definite volume but not a definite
shape
• The arrangement of atoms is more
random.
Section 3.1 Solids, Liquids, and
Gases
• LIQUIDS
Have a definite volume but no fixed shape.
The particles in a liquid are free to flow around each other
Section 3.1 Solids, Liquids, and
Gases
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• Explaining the Behavior of Liquids
Liquid particles have kinetic energy.
Particles in a liquid are more closely packed
than gas particles
A liquid takes the shape of its container because
particles in a liquid can flow to new locations.
Particles’ attractions do affect their movement.
(moving through a crowd of people)
Section 3.1 Solids, Liquids, and
Gases
• Gases
• Def.-the state of matter in which a material
has neither a definite shape nor a definite
volume
• **A gas takes the shape and volume of its
container.
Section 3.1 Solids, Liquids, and
Gases
• GASES
Have neither definite or fixed shape or volume.
The particles in a gas are: widely disbursed,
interact weakly,
move independently at high speed,
and completely fill any container they occupy.
Section 3.1 Solids, Liquids, and
Gases
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• Explaining the Behavior of Gases: Motion in
Gases
Particles in a gas are never at rest.
Some particles move faster or slower than avg.
speed of gas particles
Particles can collide with container walls or each
other causing one atom to slow down (lose
kinetic energy) or speed up (gain kinetic
energy).
The constant motion of particles in a gas allows
a gas to fill a container of any shape or size.
Section 3.1 Solids, Liquids, and
Gases
• Other States of Matter
• Most matter is solid, liquid, or gas.
• 99% of all matter observed in the universe
exists in a state not common to Earth
• Extremely high temps (sun, stars); matter
exists as plasma
– Gases whose particles are so hot they have
acquired an electrical charge.
Section 3.3 Phase Changes
• Characteristics of Phase Changes
• What happens when a substance changes
from one state to another?
• Phase- when at least (2) states of the
same substance are present, what each
different state is called
• Ex. Iceberg (solid) in the ocean (liquid)
Section 3.3 Phase Changes
• Characteristics of Phase Changes
• Phase change-the reversible physical
change that occurs when a substance
changes from one state to another
• Key Concept: Melting, freezing,
vaporization, condensation, sublimation,
and deposition are six common phase
changes.
Section 3.3 Phase Changes
• Temperature and Phase Changes
• Measuring the temp. of a substance as it
is heated or cooled is one way to
recognize a phase change.
• Key Concept: ********The temperature of
a substance does not change during a
phase change.
Section 3.3 Phase Changes
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• Energy and Phase Changes
Key Concept: Energy is either absorbed
or released during a phase change.
Endothermic change -the system absorbs
energy from its surroundings
The amt. of energy absorbed depends on
the substance.
**Heat of fusion (melting)-the amount of
energy absorbed for water (ice) to melt
Section 3.3 Phase Changes
• Energy and Phase Changes
• Ex. 1g of ice absorbs 334 joules (J) of
energy as it melts; 1g of water releases
334 J of energy as it freezes
• Exothermic change-the system releases
energy to its surroundings (freezing)
Section 3.3 Phase Changes
• Melting and Freezing
• Key Concept: The arrangement of molecules in
water becomes less orderly as water melts and
more orderly as water freezes.
• Melting-some molecules of water gain enough
energy to overcome the attractions and move
from their fixed positions
• When all molecules have enough energy to
move, melting is complete.
Section 3.3 Phase Changes
• Melting and Freezing
• Freezing-the avg. kinetic energy of water
molecules decreases (they move more
slowly)
• Some molecules move slow enough for
the attraction b/t molecules to have an
effect
• When all the molecules are drawn into an
orderly arrangement, freezing is complete.
Section 3.3 Phase Changes
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• Vaporization and Condensation
Vaporization-the phase change in which a
substance changes from liquid into a gas
Is an endothermic process
1g of water gains 2261J of energy when
vaporized (amt. of energy called heat of
vaporization)
2 Vaporization processes: evaporation
and boiling
Section 3.3 Phases Changes
• Vaporization and Condensation
• Key Concept: Evaporation takes place at the
surface of a liquid and occurs at temperatures
below the boiling point.
• Evaporation-the process that changes a
substance from a liquid to a gas at temps. below
the substance’s boiling point.
• Molecules near the surface move fast enough to
escape the liquid (water vapor/open container)
Section 3.3 Phase Changes
• Vaporization and Condensation
• In a closed container- water vapor collects
above the liquid
• Vapor pressure-pressure caused by the
collisions of the vapor with the container’s
walls (vapor press. ↑ as temp ↑)
• Boiling-when vapor pressure=the
atmospheric pressure (i.e.. The temp. that
this occurs at is the boiling point.)
Section 3.3 Phase Changes
• Vaporization and Condensation
• Boiling- temp. goes up; molecules speed up;
(100oC) molecules below the surface have
enough kinetic energy to overcome the
attraction of neighboring molecules
• Boiling point depends upon atmospheric
pressure.
• High elevation=atmospheric pressure is lower;
water can boil at temperatures lower than 100oC
(takes food longer to cook)
Section 3.3 Phase Changes
• Vaporization and Condensation
• Condensation- the phase change in which
a substance changes from a gas or vapor
to a liquid
• Is an exothermic process
• Ex. Dew on grass
Section 3.3 Phase Changes
• Sublimation and Deposition
• Sublimation- phase change in which a
substance changes from a solid to a gas
or vapor without changing to a liquid first
• Is an endothermic process
• Ex. Dry ice; cold CO2 vapor causes water
vapor in the air to condense and form
clouds
Section 3.3 Phase Changes
• Sublimation and Deposition
• Deposition- when a gas or vapor changes
directly into a solid without changing into a liquid
first
• Is an exothermic process and the reverse of
sublimation
• Ex. Frost on windows; water vapor in air makes
contact w/ cold glass; it loses enough kinetic
energy to change from gas to solid
Section 3.3 Phase Changes
• Endothermic Phase Changes (energy is
absorbed; state of matter goes from more
orderly to less orderly arrangement of
particles)
• Melting
• Vaporization
• Sublimation
Section 3.3 Phase Changes
• Exothermic Phase Changes (energy is
released; states of matter go from less
orderly to more orderly arrangement of
particles)
• Freezing
• Condensation
• Deposition
Figure 16
Phase Changes