Gas
Liquid
Solid
Highly Compressible
Slightly Compressible
Very slightly compressible
Low Density
High Density
High Density
Fills container completely
Does not expand to fill
container
Rigidly retains its volume
Assumes shape of
container
Assumes shape of
container
Retains its own shape
Rapid diffusion
Slow diffusion
Extremely slow diffusion,
surface only
High expansion on heating Low expansion on heating
Low expansion on heating
Total Disorder
Ordered arrangement
Disordered
 Forces of attraction between neighboring
particles
 Much weaker than bonding forces
 Responsible for state of matter and some
physical properties
 e.g., The stronger the attractive forces, the higher the
melting and boiling points
 Also involved in change of state
 London Dispersion forces
 Dipole-dipole forces
 Hydrogen bonds
 The motion of electrons can create an
instantaneous dipole moment on an atom
 For example, if at any one time both of a helium
atom’s electrons are on the same side of the atom at
the same time
 A temporary dipole on one atom can cause, or
induce, a temporary dipole on an adjacent atom
 These forces are significant only when
molecules are very close together, as in a
compressed gas
 These forces are found only in nonpolar
compounds
 Molecules and atoms will lose their spherical
shape
•More compact
molecules have smaller
surface areas, weaker
London dispersion
forces, and lower
boiling points.
•Flatter, less compact
molecules have larger
surface areas, stronger
London dispersion
forces, and higher
boiling points.
 Polar molecules have a positive end and a
negative end
 Dipole-dipole forces occur when the positive end
of one molecule is attracted to the negative end
of another
 Only effective when polar molecules are very
close together
 For molecules of about the same size, dipole
forces increase with increasing polarity
If two neutral
molecules, each
having a permanent
dipole moment,
come together such
that their oppositely
charged ends align,
they will be
attracted to each
other.
 Type of dipole-dipole force
 Not a true bond!
 Occurs between molecules containing a
hydrogen atom bonded to a small, highly
electronegative atom with at least one lone pair
of electrons (e.g., N, O & F)
 The hydrogen in one molecule will be attracted
to the electronegative atom in another molecule
 Hydrogen has no inner core of electrons, so a
dipole will expose its concentrated charge on the
proton, its nucleus.
 Hydrogen can approach an electronegative
atom very closely and interact strongly with it.
•Electron shell around a hydrogen
atom is rather thin, giving the
hydrogen atom a small positive
charge.
•Electron shell round an oxygen atom
is quite thick, and so oxygen carries an
extra bit of negative charge.
•These opposite charges attract,
although quite weakly.
•This weak force is called a hydrogen
bond. The hydrogen atoms of one
water molecule stick to the oxygen
atoms of nearby water molecules.
 Have much greater densities than their vapors
 Only slightly compressible; not a discernable
difference when compressed
 Fluidity: ability to flow
 Liquids can diffuse through one another, but at a
much slower rate than gases
 Viscosity: resistance to flow
 Determined by the type of intermolecular forces
involved, the shape of the particle, and the
temperature
 The stronger the attractive forces, the higher the
viscosity
 The larger the particles, the higher the viscosity
 Increases as temp decreases
 Surface Tension: the imbalance of forces at the
surface of a liquid
 The uneven forces make the surface behave as
if it has a tight film stretched across it
 The stronger the intermolecular forces, the
higher the surface tension

Surfactants: compounds that lower the
surface tension of water


Frequently added to detergents
Capillary action: movement of a liquid through
narrow spaces
 Have extremely strong intermolecular forces in
order for solids to have definite shape and
volume
 Particle arrangement causes solids to almost
always have higher densities than liquids
 Ice is an exception: it expands when it freezes
because of the way the particles arrange themselves
during the freezing process
 Particle arrangements cause different types of
solids:
 Crystalline solids




Molecular solids
Covalent network solids
Ionic solids
Metallic solids
 Amorphous solids
 Has atoms, ions, or molecules arranged in an
orderly, geometric, 3-D structure
 Individual pieces of a crystalline solid are called
crystals
 Smallest arrangement of connected points that
can be repeated in 3 directions to form a lattice
is called a unit cell
 There are 7 different crystal systems based on
shape
 Held together by dispersion forces, dipole-dipole
forces or hydrogen bonds
 NOT held together by genuine bonds (ionic and
covalent)
 Most are NOT solids at room temperature
 Poor conductors of heat and electricity (don’t
contain ions)
 Examples are sucrose and ice
–Molecular such as sucrose or ice whose
constituent particles are molecules held together
by the intermolecular forces.
Arrangement of molecules in
liquid water
Arrangement of molecules
in ice
 Atoms that can form multiple covalent bonds
 Form a network of atoms that do not have a unit
cell
 Most allotropes exist in this form
 Allotropes are forms of the same element that have
different bonding patterns of arrangement
 Examples include diamonds and graphite,
quartz
Graphite
Diamond
Covalent network solids such as quartz where
atoms are held together by 3-D networks of
covalent bonds. Here the hexagonal pattern of Si
(violet) and O (red) atoms in structure matches
the hexagonal crystal shape
 Type of crystalline solid
 Type and ratio of ions determine the structure of
the lattice and the shape of the structure
 The network of attractions that extend through
an ionic compound gives these compounds their
high melting points and hardness
 Strong but brittle
 When struck, cations and anions are shifted,
which causes repulsion that in turn shatter the
crystal
 Poor conductors of heat and electricity in solid
form
•Ionic solids are an orderly pattern of one ion,
generally the anion, with cations positioned in
'holes' between the anions
•The occupation of these 'holes' depends on
the formula of the ionic compound
Sodium chloride
Cupric chloride
 Consist of positive metal ions surrounded by a
sea of mobile electrons
 Mobile electrons make metals malleable, ductile,
and good conductors of heat and electricity
•A series of metals atoms that
have all donated their valence
electrons to an electron cloud
that permeates the structure
•This electron cloud is
referred to as an electron sea
•Visualize the electron sea
model as if it were a box of
marbles that are surrounded
by water. The marbles are the
metal atoms and the water
represents the electron sea.
•The marbles can be
pushed anywhere within
the box and the water
will follow them, always
surrounding the
marbles.
•This unique property,
allows metallic bonds to
be maintained when
pushed and pulled in all
sorts of ways.
•As a result, they are
malleable and ductile.
Gold
Copper
Silver
 Solid in which the
particles are not
arranged in a regular,
repeating pattern, but
still retain rigidity
 Examples include
glass, rubber, many
plastics, tar and wax
 Particles are trapped
in a disordered
arrangement that is
characteristic of
liquids
 Always involve
a change in
energy
 Energy is
needed either
to overcome or
form attractive
forces between
particles
 Melting point/freezing point: temp at which solid
and liquid forms exist in equilibrium
 Melting is endothermic
 Freezing is exothermic
 The change of state from a liquid to a gas
 Endothermic process
 Two methods of vaporization:
 Evaporation
 Boiling
 Occurs at the surface of a liquid
 Occurs because molecules close to the surface
have enough energy to overcome the attractions
of neighboring molecules and escape
 Slower molecules stay in the liquid state
 Rate of evaporation increases as temp
increases
 Occurs within the liquid
 Boiling point: temp at which vapor pressure
equals atmospheric pressure
 If vapor pressure is less than atmospheric
pressure, bubbles do not form
 Change of a gas to a liquid
 Exothermic process
 Molecules of vapor can return to the liquid state
by colliding with the liquid surface
 The particles become trapped by the
intermolecular attractions of the liquid
 Sublimation: solid goes directly to a gas without
passing through the liquid phase
 Deposition is the reverse process
 Sublimation is endothermic
 Deposition is exothermic
 Graphic illustrations of phase changes
 Plot of temp of a sample as a function of time
 Notice temp remains constant during phase
changes while amount of energy varies
Heating Curve of Water
A: Rise in temperature as ice absorbs heat.
B: Absorption of heat of fusion.
C: Rise in temperature as liquid water absorbs heat.
D: Water boils and absorbs heat of vaporization.
E: Steam absorbs heat and thus increases its temperature.
The above is an example of a heating curve. One could reverse the process, and
obtain a cooling curve. The flat portions of such curves indicate the phase changes.
 Diagram that relates the states of a substance to
temp and pressure
 State depends on temp and pressure
 2 states can exist simultaneously at certain
temps and pressures
 Triple point: the temp and pressure when all
three states exist at the same time
•TRIPLE POINT - The temperature and pressure at which the solid, liquid, and
gas phases exist simultaneously.
•CRITICAL POINT - The temperature above which a substance will always be
a gas regardless of the pressure.
•FREEZING POINT - The temperature at which the solid and liquid phases of
a substance are in equilibrium at atmospheric pressure.
•BOILING POINT - The temperature at which the vapor pressure of a liquid is
equal to the pressure on the liquid.
•Normal (Standard) Boiling Point - The temperature at which the
vapor pressure of a liquid is equal to standard pressure (1.00 atm = 760
mmHg = 760 torr = 101.325 kPa)
•NOTE –
•The line between the solid and liquid phases is a curve of all
the freezing/melting points of the substance.
•The line between the liquid and gas phases is a curve of all the
boiling points of the substance.