Chapter 13: States of Matter1
Section 13.1 – Gases
Types of atoms present (composition) and
their arrangement (structure) determine
the chemical properties of matter.
Kinetic-Molecular Theory
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Describes the behavior of gases in terms
of particles in motion in terms of size,
motion, and energy
All matter is in constant motion
o Size: small particles that are
separated by empty space (volume of
particles is small compared to the
volume of empty space)
 No significant attractive or
repulsive forces between them
o Motion: constant, random motion;
particles move in a straight line until
they collide with other particles or
with the walls of the container
 Perfectly elastic collisions – hit
and bounce off, without loss of
kinetic energy
o Energy: kinetic energy of a particle
is determined by its mass and its
velocity. KE = ½mv2.
 All particles have the same
mass, but not the same
velocity
 All particles do not have the
same kinetic energy.
 Kinetic energy is related to
temperature.
 Temperature is the average KE
of the particles in a sample of
matter.
 Change in temperature
indicates a change in
velocity since mass is
constant
 Absolute zero (0 K), all
molecular motion stops
Kinetic-molecular theory explains the
behavior of gases
The constant motion of gas particles
allows a gas to expand until it fills its
container.
Properties of gases
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Low density: density = mass/volume
Gases have low density because the
gas particles are separated by a lot
of space.
Compression and expansion: large
amount of empty space between gas
particles allow the gas to be
compressed (squeezed into a smaller
space). When allowed, the gas
particles will expand to fill all
available space.
Diffusion and effusion
o No attraction or repulsion
between gas particles
o Gas particles flow past each
other
o Random motion of the gas
particles causes the gases to
mix until they are evenly
distributed (diffusion).
 Diffusion always occurs
from high concentration
to low concentration
 Light particles diffuse
faster than heavier
particles
o Effusion: a gas escapes
through tiny openings (balloon
deflates over time, puncture
in a tire)
Absolute zero = -273°C
Chapter 13: States of Matter2
Graham’s Law of Effusion
Gas Pressure
Rate of effusion for a gas is inversely
proportional to the square root of its
molar mass
Pressure is defined as force per unit
area. P = F
A
Rate of effusion ∝
𝟏
𝒎𝒐𝒍𝒂𝒓𝒎𝒂𝒔𝒔
Gas particles exert pressure when they
collide with the walls of their container.
Also applies to rate of diffusion:
𝑹𝒂𝒕𝒆𝑨
𝑹𝒂𝒕𝒆𝑩
=

𝒎𝒐𝒍𝒂𝒓 𝒎𝒂𝒔𝒔 𝑩
𝒎𝒐𝒍𝒂𝒓 𝒎𝒂𝒔𝒔 𝑨

Compare the diffusion rate of ammonia
(NH3) and hydrogen chloride (HCl).

Ammonia has a molar mass of 17.0 g/mol;
hydrogen chloride has a molar mass of
36.5 g/mol. What is the ratio of their
diffusion rates?
Write down the givens:
Molar mass NH3 = 17.0 g/mol
Molar mass HCl = 36.5 g/mol
An atmosphere surrounds the Earth that
extends into space. Those air particles
exert pressure in all directions.
Substitute into Graham’s rate of
diffusion equation:
Rate NH3 =

𝒎𝒐𝒍𝒂𝒓 𝒎𝒂𝒔𝒔 𝒐𝒇 𝑯𝑪𝒍

𝒎𝒐𝒍𝒂𝒓 𝒎𝒂𝒔𝒔 𝒐𝒇 𝑵𝑯𝟑
Rate HCl
=
𝟑𝟔.𝟓 𝒈/𝒎𝒐𝒍
𝟏𝟕 𝒈/𝒎𝒐𝒍

= 𝟐. 𝟏𝟓 = 1.47
Ammonia molecules diffuse about 1.5
times faster than hydrogen chloride
molecules.
Individual gas particles have little
mass, so will have little pressure
When 1022 particles of gas are in a
liter container, they can exert
significant pressure.
SI unit of pressure – Pascal (Pa)
which equals 1 Newton of force
per square meter (N/m2)
o Pascal is a small unit, so
kilopascal (kPa) is used more
often
Atmospheric pressure = pressure
exerted by a 1-kilogram mass on a
square centimeter.
Air pressure varies at different
points on Earth (less on mountain
tops than at sea level)
1.01 x 105 Pa or 101.325 kPa
Measuring air pressure


Barometer (closed tube face down
into a liquid, usually mercury
o Difference in height of fluids in
tubes indicates pressure
Manometer – flask is connected to a Utube that is connected to a vacuum.
o Fluid is equal in arms when the
vacuum is closed
o When gas is released, the
height is no longer equal
Chapter 13: States of Matter3
Dalton’s law of partial pressures
Each gas in a mixture exerts pressure independently of the other gases present.
Dalton’s law of partial pressures: total pressure of a mixture of gases is equal to the
sum of the pressures of all the gases in the mixture.
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Portion of the total pressure contributed by a single gas is called its partial
pressure.
PTotal = P1 + P2 + P3 … PN
Partial pressure of a gas depends on
o Moles of the gas
o Size of the container
o Temperature of the gas
Section 13.2 – Forces of Attraction
2 types of forces:


Intramolecular forces
o Forces within an atom
o Ionic, covalent, and
metallic bonds
Intermolecular forces
o Forces between atoms
o Weaker than the
intramolecular (bonding)
forces
Dispersion Forces
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3 types of intermolecular forces:

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Dispersion forces
Dipole-dipole forces
Hydrogen bonding

Weak force created by the temporary
shifts in the density of electrons in
electron clouds.
Electrons in the electron clouds are in
constant motion; occasionally more
electrons will surround a particular
nucleus, giving it a greater density of
electrons.
Not permanent
Only important when no other forces
are present
o Dominant force between
identical nonpolar molecules (ex.
O2)
Force increases as the number of
electrons involved increases
o Halogens: F, Cl, Br, I – number of
nonvalence electrons increases
as you move down the periodic
table
o F and Cl are gases at room
temperature
o Br is a liquid at room
temperature
o I is a solid at room temperature
Chapter 13: States of Matter4
Dipole-dipole forces

Attractions between oppositely
charged regions of polar molecules
o Neighboring molecules align
themselves positive to negative
o Permanent
o Stronger than dispersion as
long as the molecules being
compared are approximately
the same mass.
Hydrogen Bonds
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
Section 13.3 – Liquids and Solids
Kinetic-molecular theory applies to both
liquids and solids, not just gases.

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Consider the forces of attraction
between the particles
Consider their energy of motion
Liquids
Properties:

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No definite shape; particles are able
to slide past each other
Definite volume
Forces between particles are stronger
than in gases, so liquids are more
dense than gases
Have fluidity – ability to flow
Viscosity – measure of the resistance
of a liquid to flow
o Determined by:
o Intermolecular force
stronger the force, the higher
the viscosity
o Shape of the particles
larger molecules flow slower
than smaller molecules
o Temperature
lower temperatures increases
viscosity

Hydrogen bonds are one type of dipoledipole attraction
Occurs between molecules containing a
hydrogen atom bonded to a small,
highly electronegative atom, with one
lone electron pair
o Fluorine, oxygen, or nitrogen
o Because of the electronegativity,
hydrogen acts almost like a bare
proton
o Hydrogen is only element that
can do this because it lacks inner
electrons
Hydrogen end of a molecule is strongly
attracted to the negative end of other
molecules
Effects of hydrogen bonds
1. Increase in melting and boiling points
2. Bonds with hydrogen are stronger than
other compounds with the same
electronegativity difference
Additional properties of liquids


Surface tension – measure of the
inward pull by particles in the
interior of the liquid
o Intermolecular forces are not
equal on all particles in a liquid
o Particles at the surface have no
attraction from above, so they
adhere more strongly to the
particles below
o Particles tend to occupy the
smallest area which creates
tension
o Stronger the attraction, the
greater the surface tension
Surfactants: compounds that lower
the surface tension of water
o Soaps, detergents
Chapter 13: States of Matter5
Liquid properties, continued

Capillary action – liquid rising through a tube due to attractive forces
o Water rises until the forces between the water, the surface and gravity are equal
o Adhesion – attractive forces between molecules that are different
o Cohesion – attractive forces between molecules that are identical
Solids
Types of solids:
Properties:
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Definite shape
Definite volume
Strongest attractive forces between
particles
Most dense
Crystalline
Molecular
Covalent network
Ionic
Metallic
Liquid crystals
Amorphous
Crystalline solids
Crystalline – rigid body in which particles are arranged in a regular repeating pattern

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Arrangement is determined by the bonds between the crystals
Have well-defined melting points
Have sharp edges
o Unit cell – smallest repeating unit of a crystalline solid
Chapter 13: States of Matter6
Types of Crystalline Solids – Classified into 5 categories based on particle types:

Atomic solids
o Atoms
o Soft to very soft
o Very low melting points
o Poor conductivity
o group 8A (noble gases)

Molecular solids
o Molecules
o Fairly soft; not normally solid at room temperature
o Low to moderately high melting points
o Poor conductivity of heat and electricity
o Held together by dispersion forces, dipole-dipole forces, or hydrogen bonds
o I2, H2O, NH3, CO2, table sugar (solid at room temp due to large molar mass)

Covalent network solids
o Atoms connected by covalent bonds; form multiple covalent bonds
o Very hard
o Very high melting points
o Often poor conductivity
o Diamond (C), quartz (SiO2)

Ionic
o
o
o
o
o

Metallic solids
o Nucleic surrounded by shared electrons
o Soft to hard
o Low to very high melting points
o Malleable and ductile
o Excellent conductivity (heat and electricity)
o All metallic elements
solids
Ions surrounded by oppositely charged ions
Very hard, brittle
High melting points
Poor conductivity
NaCl, KBr, CaCO3
Amorphous solids
Solids in which the particles are not
arranged in a regular, repeating pattern

“without shape”

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Forms when a molten material cools
too quickly to allow time for crystals to
form
Glass, rubber, plastic, wax
Over long periods of time, the material
will flow
Chapter 13: States of Matter7
Section 13.4 – Phase Changes
Terms
Gas – substances that are gaseous at room
temperature
Vapor – substances that are in the gaseous
state but are solids or liquids at room
temperature
Vapor equilibrium – number of molecules
leaving the liquid state equals the number
returning to the liquid state in a closed
container
Dynamic equilibrium – constant motion of
particles
When energy is added or removed from a
system, one phase can change into another.
Phase Changes require energy
Melting (solid to liquid)
Imagine a glass containing ice and water
Heat is the transfer of energy from an object
at a higher temperature to an object at a
lower temperature
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Vapor pressure – pressure exerted by a
vapor in equilibrium with its liquid
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Substances with low vapor
pressures have strong
intermoleulcar forces and viceversa
Ionic compounds do not exert any
significant vapor pressure
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The addition of energy disrupts the
hydrogen bonds holding the water
molecules together in the ice crystal
Enough energy is added, the bonds
break, molecules move apart and
enter the liquid phase
o Energy is used to break the
hydrogen bonds, not raise the
temperature of the ice
When temperature is raised, kinetic
theory tells us that the velocity of the
particles will increase (KE = ½mv2)
o More collisions between the
particles
o Order of the solid is lost
Vapor pressure of the solid and vapor
pressure of the liquid are equal
Non-polar molecules have lower
melting points than polar molecules of
the same mass.
Ionic bonds are much stronger than
hydrogen (or covalent bonds). More
energy is required to melt ionic
substances.
Melting point of a crystalline solid:
temperature at which the forces holding the
crystal lattice are broken and it becomes a
liquid.
Chapter 13: States of Matter8
Vaporization (liquid to gas)
Imagine a glass containing ice and water
While the ice melts, the temperature of
the ice-water mixture remains constant.
After melting …

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Additional energy increases the KE of
the liquid molecules
When KE is great enough, particles
enter the gas phase and bubbles form
Evaporation – vaporization that only occurs
at the surface

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Boiling takes place throughout the
liquid
Boiling point: temperature at which
the vapor pressure of a liquid equal
the external or atmospheric
pressure
o Boiling point decreases as
altitude increases
o Evaporation only occurs at the
surface, boiling takes place
throughout the liquid
Phase diagrams
Graph of temperature vs. pressure that
shows in which phase a substance exists at
various temperatures and pressures
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At any point on a line, a substance can
exist in two states
Triple point – point where all 3 phases
can coexist
o All 6 phase changes can occur
Critical point – point at which a
substance cannot exist as a liquid
Sublimation (solid to gas, without going
through the liquid phase)
Solids have high enough vapor pressure at
room temperature to vaporize rapidly if not
kept contained
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Solid CO2 (dry ice)
Solid iodine
Air fresheners
Ice cubes left in the freezer for a long
time
Phase Changes that Release Energy
Deposition (gas to a solid, without going
through the liquid phase)
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Reverse of sublimation
Frost
Snow
Condensation (gas to liquid)
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Formation of hydrogen bonds; vapor
phase changes to the liquid phase
Reverse of vaporization
Condensation point and boiling point
are the same
Liquefaction – condensing substances
that are gaseous at room temperature
Critical temperature: temperature at
which a gas will not condense
regardless of pressure
o Indicates the strength of the
forces between the molecules
Freezing (liquid to solid)
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Reverses of melting
Particles lose energy until they no
longer can slide past one another
Particles settle into an orderly
arrangement