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Chapter 13:
States of Matter
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13. 1 Gases
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Kinetic Molecular Theory—attempts to explain the
properties of gases.
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Assumes:

Particles are small and separated by empty space.

No significant attractive or repulsive forces.

Particles are in constant, random motion.

Collisions are elastic (no kinetic energy lost)
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Kinetic energy determined by KE = ½ mv2.
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Temperature
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Although all particles of the same gas have the same mass,
they don’t have the same velocity or kinetic energy.
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Temperature is a measure of the average kinetic energy of
the particles.

At a given temperature, all gases have the same kinetic
energy.
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Behavior of Gases
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KM theory explains the behavior of gases.
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Low density: Due to the large amount of space between the
particles in the gas.
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Compression and expansion: Gases expand to fill their
containers. Becomes more dense in a smaller container; less
dense in a larger container.
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Diffusion: As gas particles flow past one another, lighter particles
travel more quickly (higher velocity).
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Effusion: The escape of a gas through a tiny hole.
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Graham’s Law of Effusion
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States that the rate of effusion for a gas is inversely
proportional to the square root of its mass.
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To compare the rates of diffusion of two gases use the
formula:
RateA = sqrt (molar massB)
RateB
(molar massA)
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Gas Pressure
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Occurs because of the collision of gas particles with the walls
of their container.

Pressure is defined as force per unit area.

The particles in Earth’s atmosphere exert pressure on Earth
known as atmospheric pressure, or air pressure.

At higher elevations, air pressure is lower than at sea level.
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Pressure
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Air pressure is measured using a barometer or a
manometer.

The SI unit for pressure is the pascal (Pa).
1 Pa = 1 N / 1 meter2
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There are other units to measure pressure, which include
atmosphere (atm), torr, bar, pounds per square inch (psi) and
kilopascals (kPa).
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Pressure Unit Relationships
Unit
Relationship To Other Units:
Atmosphere (atm)
1 atm = 760 mmHg = 760 torr = 760 bar= 14.7
psi
Pascal (Pa)
0.1013 Pa = 1 atm
millimeters Mercury
(mm Hg)
760 mm Hg = 1 atm = 760 torr =
Torr (torr)
1 torr = 1 mm Hg = 1.33 X 10-3 barr
Kilopascals (kPa)
1 atm = 101.3 kPa
Pounds per square
inch (psi)
1 psi = 51.71 torr
Bar (bar)
1 bar = 1 X 102 kPa = 0.98692 atm
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Dalton’s Law of Partial Pressures
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Dalton’s law of partial pressures states that the total
pressure of a mixture of gases is equal to the sum of the
pressures of all the gases in the mixture.
Ptotal = P1 + P2 + P3 + …Pn
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The partial pressure of a gas depends on the number of
moles of the gas, the size of the container, and the
temperature of the mixture.
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Collection of Gas Over Water
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Dalton’s law can be used to determine the amount of gas
produced by a reaction.

When the gas is collected in an inverted container over
water, the gas will displace the water.
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If the gas doesn’t react with water, the total pressure will be a
combination of the gas and the water vapor, which has a fixed
value.
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Forces of Attraction
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The forces that hold particles together in ionic, covalent and
metallic bonds are called intramolecular forces.

Intermolecular forces are weaker than intramolecular
forces.
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There are three types of intermolecular forces: dispersion
forces, dipole-dipole forces, and hydrogen bonds.
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Dispersion Forces
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These forces are also called London forces.
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They are weak forces that result from temporary shifts in the
density of electrons in electron clouds.
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When two nonpolar molecules are close to one another the
electron cloud in one molecule will repel the the electron
cloud of the other. This causes the electron density to shift,
creating a temporary dipole.
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Dispersion Forces (cont’d)
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Weak forces of attraction form between oppositely charged
regions.

These forces are only important when no stronger force is
acting.
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They are noticeable between identical nonpolar molecules.
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Dipole-Dipole Forces
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Attractions between oppositely charged regions of polar
molecules are called dipole-dipole forces.
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Polar molecules have permanent dipoles and orient so that
opposite charges are attracted.
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These forces are stronger than dispersion forces.
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Hydrogen Bonds
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A hydrogen bond is a special type of dipole-dipole
attraction that occurs between molecules containing
hydrogen and either fluorine, oxygen or nitrogen.
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Liquids and Solids
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The KM theory also explains the behavior of liquids and
solids; however they exhibit much stronger forces of
attraction.
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Liquids
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They take the shape of their container but have a fixed
volume. Their forces of attraction limit their range of motion.
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Density: Higher in liquids because particles are held
together.
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Compression: requires enormous pressure, produces small
volume change.
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Fluidity: less fluid than gases.
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Liquids (cont’d)
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Viscosity: stronger the intermolecular forces, the higher the
viscosity; increases with mass; decrease with temperature
increase.
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Surface tension: measure of the inward pull by particles to
the interior of a liquid. Stronger attractions means higher
surface tension.
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Capillary action: movement up a narrow glass tube. Occurs
when adhesion (intermolecular) is greater than cohesion
(intramolecular).
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Solids
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Strong attractive forces that limit movement to vibrations.
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Definite shape and volume.
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Particles are more closely packed than in liquids, therefore a
higher density.
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Normally, liquids are less dense than solids, though water is
an exception.
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Crystalline Solids
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A solid whose atoms, ions, or molecules are arranged in an
orderly, geometric, three dimensional structure (lattice) are
called crystalline solids.
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Not all solids are crystalline. In amorphous solids there is
no regular, repeating pattern.
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Phase Changes
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Most substances can exist in three states—solid, liquid, gas—
depending on their temperature and pressure.
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These states can coexist; however, if energy is added or
taken away, one phase can change into another.
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Phase Changes That Require Energy
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The three phases that require energy are melting,
vaporization and sublimation.
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The temperature at which the forces holding the crystal
lattice together are broken and a solid becomes a liquid is
known as the melting point.
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The process by which a liquid becomes a gas or vapor is
known as vaporization.
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Vaporization on the surface is called evaporation.
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The temperature at which liquids become vapors is called
the boiling point.
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Sublimation
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Sublimation is the process in which a solid changes directly
to a gas without first becoming a liquid.
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Examples of materials that sublime include solid air
fresheners, dry ice, moth balls, etc.
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Phase Changes That Release Energy
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There are three phase changes that release energy:
condensation, freezing and deposition.
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Condensation is the process of a gas or vapor becoming a
liquid.
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The freezing point of a material is the temperature at which a
liquid becomes a solid.
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Going from a gas to a solid, without becoming a liquid is
known as deposition.
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Six Phase Changes
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Phase Diagrams
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Temperature and pressure control the phase of a substance.
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A phase diagram is a graph of pressure versus temperature
that shoed in which phase a substance exists at different
temperatures and pressures.
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There are typically three regions and three curves. The
curves represent where two phases coexist.
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Triple Point
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The triple point is the point on a phase diagram that
represents the temperature and pressure at which the three
phases of a substance can coexist. All 6 phases occur here.
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The critical point shows the critical temperature and
pressure, above which a substance cannot exist as a liquid.
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Phase Diagram