Chapter 10
States of Matter
10.1 The Kinetic-Molecular
Theory of Matter
Kinetic-Molecular Theory of Gases
• Particles of matter are ALWAYS in motion;
constant, rapid motion. (kinetic energy!)
• Particles are very small & relatively far
apart.
• Collisions of particles with container walls
cause pressure exerted by gas.
• Volume of individual particles is  zero.
• Particles exert no attractive or repulsive
forces on each other.
Kinetic-Molecular Theory of Gases
• Gas particles undergo elastic collisions:
Collisions in which no energy is lost
Air Hockey Table
• Average kinetic energy is directly proportional
to Kelvin temperature of a gas.
Ideal Gas
• An imaginary gas that perfectly fits all the
assumptions of the kinetic-molecular theory
• A gas with its particles in constant random
motion without attraction for each other is
called an Ideal Gas. These particles undergo
elastic collisions.
• Nearly all real gases behave as ideal gases
EXCEPT at very low temperatures or high
pressures.
Real Gases
• A gas that does not behave completely
according to the assumptions of the kineticmolecular theory.
• Real gases occupy space and exert attractive
forces on one another
Likely to behave nearly
ideally:
Gases @ high temp. & low
pressure
Small non-polar gas molecules
Likely not to behave ideally:
Gases @ low temp. & high
pressure
Large polar gas molecules
Kinetic-Molecular Theory of the
Nature of Gases
• Expansion
Gases do not have a definite shape or volume
Gases take the shape of their containers
Gases evenly distribute themselves within a
container
• Fluidity
Gas particles easily flow past one another
• Low Density
A substance in the gaseous state has 1/1000 the density of the
same substance in the liquid or solid state
• Compressibility
Gases can be compressed, decreasing the distance between
particles, and decreasing the volume occupied by the gas
Kinetic-Molecular Theory of the
Nature of Gases
• Diffusion
Spontaneous mixing of particles
of two substances caused by
their random motion
– Rate of diffusion is dependent
upon:
• speed of particles
• diameter of particles
• attractive forces between the
particles
Kinetic-Molecular Theory of the
Nature of Gases
• Effusion
Process by which
particles under
pressure pass through
a tiny opening
– Rate of effusion is
dependent upon:
• speed of particles
(small molecules have
greater speed than
large molecules at the
same temperature, so
the effuse more rapidly)
Chapter 10
States of Matter
10.2 Liquids
Some Properties of a Liquid
Surface Tension:
The resistance to an
increase in its surface
area (polar molecules,
liquid metals).
A force that tends to pull
adjacent parts of a
liquid's surface together,
thereby decreasing
surface area to the
smallest possible size.
Some Properties of a Liquid
Capillary Action:
Spontaneous rising of
a liquid in a narrow
tube.
Some Properties of a Liquid
Viscosity: Resistance to
flow (molecules with
large intermolecular
forces).
Some Properties of Liquids
Volatility
• Liquids that have weak forces of
attraction and evaporate easily
Nonvolatile Liquids
• Liquids that have strong forces of
attraction and do not evaporate easily
Properties of Fluids
Relative High Density
• 10% less dense than solids (average)
• Water is an exception
• 1000x more dense than gases
Relative Incompressibility
• The volume of liquids doesn't change
appreciably when pressure is applied
Ability to Diffuse
• Liquids diffuse and mix with other liquids
• Rate of diffusion increases with temperature
Chapter 10
States of Matter
10.3 Solids
Types of Solids
Crystalline Solids:
highly regular
arrangement of their
components
[table salt (NaCl),
pyrite (FeS2)].
Types of Solids
Amorphous solids aka
supercooled liquids:
considerable disorder in
their structures (glass).
• Greek for "without
shape"
• Formation of amorphous
solids:
• Rapid cooling of molten
* They do not have definite
materials can prevent
the formation of crystals melting points
Representation of Components in
a Crystalline Solid
Lattice: A 3-dimensional
system of points
designating the centers of
components (atoms, ions,
or molecules) that make up
the substance.
Types of Crystalline Solids
Ionic Solid: contains ions
at the points of the lattice
that describe the structure
of the solid (NaCl).
Unit Cell
• The smallest portion of a crystal
lattice that shows the threedimensional pattern of the entire
lattice
Types of Crystalline
Solids
Molecular Solid: discrete
covalently bonded
molecules at each of its
lattice points (sucrose,
ice).
Packing in Metals
Model: Packing uniform, hard spheres to best use
available space. This is called closest packing.
Each atom has 12 nearest neighbors.
Closest Packing
Holes
Metal Alloys
Substitutional Alloy:
some metal atoms
replaced by others of
similar size.
• brass = Cu/Zn
Metal Alloys
(continued)
Interstitial Alloy:
Interstices (holes) in
closest packed metal
structure are occupied
by small atoms.
steel = iron + carbon
Network Solids
•Composed of strong directional covalent
bonds that are best viewed as a “giant
molecule”.
- brittle (non-flexible)
- do not conduct heat or electricity
- carbon, silicon-based
•graphite, diamond, ceramics, glass
Sulfur – S8
Phosphorus – P4
Diamond
Graphite
Zirconia
Chapter 10
States of Matter
10.4 Changes of State
Equilibrium
• Dynamic condition in which two opposing
changes occur at equal rates in a closed
system
• A closed system at constant temperature will
reach an equilibrium position at which the
rates of evaporation and condensation will be
the same
Equilibrium Vapor Pressure
• The pressure of the
vapor present at
equilibrium.
• Determined principally
by the size of the
intermolecular forces in
the liquid.
• Increases significantly
Increasing the temperature
with temperature.
will move more particles
• Volatile liquids have
into the vapor phase to
high vapor pressures.
compensate for the new
energy
Boiling
The conversion of a liquid
to a vapor within the liquid
as well as at its surface. It
occurs when the
equilibrium vapor pressure
of the liquid equals the
atmospheric pressure
Boiling Point
• The temperature at which the equilibrium vapor
pressure of the liquid equals the atmospheric
pressure
Water boils at 100 °C at 1 atm pressure
Water boils above 100 °C at higher pressures
Water boils below 100 °C at lower pressures
LeChatelier’s
Principle
When a system at
equilibrium is placed
under stress, the system
will undergo a change in
such a way as to relieve
that stress.
Translation:
When you take something away from a system
at equilibrium, the system shifts in such a way
as to replace what you’ve taken away.
When you add something to a system at
equilibrium, the system shifts in such a way
as to use up what you’ve added.
LeChatelier’s Example #1
A closed container of ice and water at
equilibrium. The temperature is raised.
Ice + Energy  Water
The equilibrium of the system shifts to
right to use up the added energy.
the _______
LeChatelier’s Example #2
A closed container of N2O4 and NO2 at
equilibrium. NO2 is added to the container.
N2O4 + Energy  2 NO2
The equilibrium of the system shifts to
left to use up the added NO2.
the _______
LeChatelier’s Example #3
A closed container of water and its vapor at
equilibrium. Vapor is removed from the system.
water + Energy  vapor
The equilibrium of the system shifts to
right to produce more vapor.
the _______
constant
Temperature remains __________
during a phase change.
Water phase changes
Phase Diagram
Represents phases as a function of temperature
and pressure.
Critical temperature: temperature above which
the vapor can not be liquefied.
Critical pressure: pressure required to liquefy
AT the critical temperature.
Critical point: critical temperature and pressure
(for water, Tc = 374°C and 218 atm).
Water
Water
Phase changes by Name
Carbon dioxide
Carbon
dioxide
Carbon
Carbon
Sulfur
Chapter 10
States of Matter
10.5 Water
Water’s Properties
Sea Ice
• Ice forms on top of the
ocean in a thin layer & acts
to insulate the warmer
waters below from the colder
air temperatures.
• This occurs in the polar
regions, the Artic & Antarctic
• Since ice is less dense than liquid water, it will float on
top, instead of sinking which would kill all life below the
surface.
•Sea ice is not the same as an iceberg. Icebergs are pieces
of glaciers which are formed by snowfall on land.
• Sea ice is not salty, as the hydrogen bonds that hold ice
together will not form properly if salt remains in the
structure.
Sea Water Density
• There are 2 main factors that affect the density of
sea water:
Temperature & Salinity
• As the temperature decreases, density increases.
• As salinity increases, density increases.
• Since the least dense layer of a liquid will float
above a more dense layer, the warmest & lowest
salinity layer will be on top.
• However, temperature has a greater effect on
density than salinity.
• That means that a higher salinity layer can float on
top of a lower salinity layer if it is considerably
warmer in temperature.
Molar Heat of Fusion
• The amount of heat energy required to melt
one mole of solid at its melting point.
• 6.009 kJ per 1 mole
Practice!
How much energy is absorbed when 16.3 g of
ice melts?
16.3 g x 1 mole x 6.009 kJ = 5.44 kJ
18 g
1 mol
Molar Heat of Vaporization
• The amount of heat energy required to vaporize
one mole of a liquid at its boiling point
• Strong attractive forces between particles result in
high molar heat of vaporization
• 40.79 kJ per 1 mole
Practice!
Find the mass of liquid water required to absorb
5.23 x 104 kJ of energy upon boiling.
5.23 x 104 kJ x 1 mole x 18 g = 2.31 x 104 g
40.79kJ 1 mol
Specific Heat
• The amount of energy required to change the
temperature of 1 gram of a substance by 1
degree Celsius is known as specific heat
capacity. (Ch.16 p.533)
Q = s * m * ΔT
Where Q = energy (heat) required
s = specific heat capacity
m = mass of sample, grams
ΔT = change in temperature, Celsius
Specific Heat Capacities of Some
Common Substances
Substance
Specific Heat Capacity (J/g°C)
Water, Liquid
4.184
Water, Ice
2.03
Water, Steam
2.0
Aluminum, s
0.89
Iron, s
0.45
Mercury, l
0.14
Carbon, s
0.71
Silver, s
0.24
Gold, s
0.13
Practice!
A piece of copper alloy with a mass of 85.0 g is
heated from 30.°C to 45°C. In the process, it
absorbs 523 J of energy as heat.
a. What is the specific heat for this copper alloy?
b. How much energy will the same sample lose if it
is cooled from 45°C to 25°C?
Answers…
a. 0.41 J/g∙K
b. 7.0 x 102 J
Specific Gravity
• A function of density.
• The ratio of density of a material to the density of
water at a specified temperature.
• The density of water is usually 1 g/cm3
• When dividing density by density, the units cancel
and therefore, specific gravity has no units!
• If the density of the material is in the units g/cm3
(or g/mL) when dividing by the density of water,
the numeric value of the material’s density is the
same, but the units cancel out.
Intermolecular Forces
(Review from Chapter 6)
Forces of attraction between different
molecules rather than bonding forces
within the same molecule.
Dipole-dipole attraction
Hydrogen bonds
Dispersion forces
Dipole-Dipole Attraction
 dipole-dipole attraction: molecules with
dipoles orient themselves so that “+” and “-”
ends of the dipoles are close to each other.
Dipole
Forces
Hydrogen Bonding
 hydrogen bonds: dipole-dipole attraction in
which hydrogen is bound to a highly
electronegative atom. (F, O, N)
Hydrogen Bonding in Water
Hydrogen Bonding in DNA
London Dispersion Forces
Dispersion forces: relatively weak forces
caused by instantaneous dipole, in which
electron distribution becomes asymmetrical.
The ONLY forces of attraction that exist among
noble gas atoms and nonpolar molecules. (Ar,
C8H18)
London Dispersion Forces
Boiling point as a measure of intermolecular
attractive forces