# Chapter 13 ```Chapter 13
States of Matter
13.1 Gases
The Kinetic-Molecular Theory
Describes the behavior of gases in terms of
particles in motion
Gas Behavior
Size
Motion
Energy
Particle Size
Particle size is small relative to the space
that surrounds them
The distance between particles is so large
that no attractive or repulsive forces exist
between gas particles
Particle Motion
 Gas particles are in constant, random motion
 Particles move in a straight line until they
collide with another particle or the walls of a
container
 Gas particle collisions are elastic. Although
particles in collision can transfer kinetic to
eachother, the total kinetic energy of the
colliding particles remains constant.
Particle Energy
All particles in a gas have the same mass,
but not the same kinetic energy.
1 2
KE  mv
2
Velocity
Mass
Temperature is the measure of the
average kinetic energy of particles in a
sample of matter
Explaining the Behavior of Gases
Low Density
Particle size is small relative to the space that surrounds them As
compared to liquids or solids, gases have much smaller densities
due to the fact that fewer particles occupy the same volume
Compression and Expansion
Particle size is small relative to the space that surrounds them
The large amount of space surrounding the particles allows
particles the room to move closer together as they are
compressed
Diffusion and Effusion
Gas particles are in constant, random motion
Gas particles tend to move from a high area of concentration to
a low area of concentration. The rate at which they diffuse is
dependent on their mass.
Graham’s Law
of Effusion
1
EffusionRate 
MolarMass
Gas escaping through a
small hole
Comparison of the diffusion rates of
two gases
RateA
MolarMass A

RateB
MolarMass B
Practice Problem
What is the ratio of effusion rates for
Nitrogen (N2) and Neon (Ne)?
By: Jordan Gaffin, Sam Bear, and Keith
Zubrow
What is Pressure?
Pressure is defined
as force applied
per unit area. We
measure air and
atmospheric
pressure with a
barometer.
Measuring Air Pressure
 Italian physicist
Evangelista Torricelli
was the first to
demonstrate that air
exerted pressure.
 He invented a device
called the barometer
that assisted him in
measuring pressure.
Gas Pressure
A manometer is an instrument used to
measure gas pressure in a closed
container.
It is a flask that is connected to a U-tube
that contains mercury.
Units Of Pressure
Unit
Compared with 1
atm
Kilopascal (kPa)
1 atm = 101.3 kPa
Millimeters of
mercury (mm Hg)
1 atm = 760 mm Hg
I kPa = 7.501 mm Hg
Torr
1 atm = 760 torr
1 kPa = 7.501 torr
Pound per square
inch (psi or lb/in)
1 atm = 14.7 psi
1 kPa = .145 psi
Atmosphere (atm)
Compared with 1
kPa
I kPa = .009869 atm
Units of Pressure
 The SI unit of pressure is the pascal (Pa).
 1 pascal = the force of 1 Newton per square
meter.
 The pressures measured by barometers and
manometers can be reported in millimeters of
mercury (mm Hg).
 There is also a unit called the torr which is equal to
1mm Hg.
 Often air pressure is reported in a unit called
atmosphere (atm).
 1 atm = 760mm Hg
Conversions
To convert the different measures of
pressure, we use factor label.
Use the given measurements:
1 atm= 101.3 kPa = 14.7 psi = 760 torr
Dalton’s Law of Partial Pressures
This law 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.
Dalton’s law of partial pressures can be
summarized as:
Ptotal = P1+P2+P3+... Pn
The P’s in Daltons Theory
Ptotal is the total pressure of a mixture of
gas.
All the other P’s (P1, P2, P3, Pn…) are the
partial pressures of each gas in the
mixture.
Example
A mixture of Oxygen, Carbon Dioxide, and
Nitrogen has a total pressure of .97 atm. What is
the partial pressure of O2 if the partial pressure
of CO2 is .70 and the partial pressure of N2 is
.12?
P.97atm = P N2 .12 atm + P CO2 .7atm + P O2 x
P O2 = .97 atm - .70 atm - .12 atm
P O2 = .15 atm
Section 13.2
Forces of Attraction
By Dorothy Raginsky, Rima Naseer, Luke Morreale,
Emil George, Greg Klein
Differences between intramolecular forces
and intermolecular forces
Dispersion Forces with examples
Dipole-dipole Forces with examples
Hydrogen Bonds with examples
Intramolecular Forces
The attractive forces that hold particles
together in ionic, covalent, and metallic
bonds.
Intra- means “within”.
Intermolecular Forces
 They hold together identical particles or two
different types of particles.
 There are three types of intermolecular forces:
Dispersion forces
Dipole-dipole forces
Hydrogen bonds
 Some intermolecular forces are stronger than
others, but all are weaker than intramolecular
forces.
Dispersion Forces
 They are weak forces that result from temporary
shifts in the density of electrons in electron
clouds.
 Dispersion forces are sometimes called London
forces after the German American physicist who
first described them, Fritz London.
 They are the weakest intermolecular force.
 They are the dominant force of attraction
between identical, nonpolar molecules.
Examples
 Fluorine, chlorine, bromine, and iodine exist as
diatomic molecules.
 Recall that the number of nonvalence electrons
increases from fluorine to chlorine to bromine to
iodine. Because the larger halogen molecules
have more electrons, there can be a greater
difference between the positive and negative
regions of their temporary dipoles, and, thus, the
stronger dispersion forces. This difference in
dispersion forces explains why chlorine and
fluorine are gases, bromine is a liquid, and
iodine is a solid at room temperature.
Dispersion Forces
Dipole-dipole Forces
 The attraction between oppositely charged
regions of polar molecules are called dipoledipole forces.
 Polar molecules contain permanent dipoles; that
is, some polar regions of a polar molecule are
always partially negative and some regions are
always partially positive.
 Neighboring polar molecules orient themselves
so that oppositely charged regions line up.
Examples
When hydrogen chloride gas molecules
approach, the partially positive hydrogen
atom in one molecule is attracted to the
partially negative chlorine atom in another
molecule.
Dipole-dipole Forces
Hydrogen Bonds
A hydrogen bond is one special type of
dipole-dipole attraction that occurs
between molecules containing a hydrogen
atom bonded to a small highly
electronegative atom with at least one lone
electron pair.
For a hydrogen bond to form, hydrogen
must be bonded to either a fluorine,
oxygen, or nitrogen atom. These atoms
are electronegative.
Examples
In a water molecule, the hydrogen atoms
have a large partial positive charge and
the oxygen atom has a large partial
negative charge. When the water
molecules approach, a hydrogen atom in
one molecule is attracted to the oxygen
atom on the other molecule.
Hydrogen Bonding
Section 13.3 Liquids
Ryan Desch, Dave Derr, Colin
Drummond,Kyle Giordano, Jake
Long, Pratt Templeton
Density and Compression
At one atmosphere of air pressure liquids
are more denser than gases.
Change in volume for liquids is much
smaller than gas because liquid particles
An enormous amount of pressure must be
applied to reduce the volume of liquid by
even a few percent.
Fluidity
 Fluidity- The ability to
flow
 A liquid is considered
as fluid because it
can flow.
 Liquids can defuse
through other liquids
due to their fluidity.
 Liquids are less fluid
than gases.
Viscosity
 Viscosity- A measure of
the resistance of a liquid
to flow
 The particles in liquid are
close enough for
attractive forces to slow
their movement.
 Viscosity decreases with
temperature (Increase in
temperature = increases
the average kinetic
energy of molecules).
 For Example- When you
pour a tablespoon of
cooking oil into a frying
pan the oil tends not to
of the pan until you heat
the oil.
Surface Tension
 Surface Tension- The
energy required to
increase the surface
area of a liquid by a
given amount.
 The stronger the
attractions between
particles, the greater
the surface tension.
 SurfactantsCompounds that
lower the surface
tension of water.
 When liquid isCapillary
placed in a
narrow container the
surface forms a concave
meniscus (The surface
dips in the center).
 This movement of the
liquid, such as water, is
called capillary action or
capillarity.
 Cohesion- The force of
attraction between
identical molecules
attraction between
molecules that are
different.
Action
Solids
Joe Bruch
Bria Collins
Melanie Spivack
Laura Betterly
Solids
Density
 Solids are more dense than liquids and gases
 All the solids remain a compact phase, where
the arrangement of molecules is very ordered
and well defined. With this in mind it is
surprising that there exists some solids (few)
less dense than liquids.
 P=m( mass)/v (volume). This is the equation
for density
Crystalline Solids
 Are arranged in fixed geometric patterns or lattices.
 Ex. ice methanol and sodium chloride
 They have an orderly arranged units and are practically
incompressible
 Crystalline solids, also, show a definite melting point
 They passed rather sharply from solid to liquid state
 There are various crystalline forms which are divided into
six crystal systems or shapes: cubic, tetragonal,
hexagonal, rhombic, monoclinic, and triclinic
Molecular Solids
 a solid composed of molecules held together by
relatively weak intermolecular forces
 are soft, and generally have low melting and boiling
temperatures
 Most solids are nonconducting when pure and are
insoluble in water, but soluble in non-polar solvents
 Ex. Sulfur, ice, sucrose, and carbon dioxide solid
Metallic Solids
 More than 80 elements in the periodic table are
metals
 High densities
 Low ionization energies
 Low electronegativities
 Usually, high deformation
 Malleable
 Ductile
 Thermal conductors
Covalent Network of Solids
-Have no discrete molecular units
-Held together by conventional covalent
bonds
-Continuous network of bonded atoms
examples – diamond and quartz
Amorphous solids
 is a solid in which there is no long-range order
in position- molecules are arranged in a
random manner
 common window glass and cotton candy and
plastic are examples
 Amorphous solids also unlike crystalline solids.
They do not have definite melting points
Phase Changes
By, Taylor Brink, Lindsay
Coloracci, Amanda Couch, Emily
Rorer, Callie Wendell, and Jeff
Wright
Endothermic Changes
(within)
Melting
 Hot flows into cold
 Heat disrupts the hydrogen bonds holding
the ice together
 These bonds are strong, therefore, lots of
energy is required
 Melting Point- temperature at which a solid
becomes a liquid- hard to determine exact
point
Vaporization vs. Evaporation
 Vaporization
 liquid changes into a gas (vapor)
 the gradual input of energy equals the amount of
molecules escaping from the surface
 Vapor pressure- pressure exerted by a vapor over a liquid
 Surface molecules are less attracted to other molecules
 Evaporation
 molecules escaping from the surface
 better in warmer temperatures because there is more
energy
 all evaporates over time depending on the amount of water
and energy
Boiling Process
 Boiling point- temperature where
vapor pressure equals the
atmospheric pressure
 At this temperature, all molecules
have enough energy to vaporize, where
as only the surface molecules could
before
Sublimation
 Transition from a solid to a gas,
skipping a liquid
 examples:
• Carbon Dioxide (dry ice)
• Air fresheners
• Ice in a freezer
 Happens when the pressure is low
Exothermic Changes
(outside)
Condensation
 Transition from gas to liquid
 Hydrogen bonds begin to form, and
the vapor looses energy
 examples
• Sides of a cold glass
• Dew
• Fog
• Rain
Deposition
 Transition from gas to solid
 Vapor settles on to a solid
Examples
• Frost
• Snowflakes
Freezing
 Transition from liquid to solid
 Water looses energy (heat) and
hydrogen bonds form, stopping the
motion of water molecules
 Freezing point- temperature where
liquid turns to solid
```