Download Condensed States of Matter: Liquids and Solids Chapter 14

Condensed States of
Matter:
Liquids and Solids
Chapter 14
Condensed States
• Liquids and solids = condensed
states because they have
significantly higher densities
than gases
Gas
Liquid
Solid
highly compressible
only slightly
compressible
only slightly
compressible
low density
high density
high density
fills container
completely
does not expand to
fill container - has
definite volume
rigidly retains its
volume
assumes shape of
container
assumes shape of
container
rapid diffusion
slow diffusion
retains own shape
extremely slow
diffusion - only
occurs at surfaces
high expansion on
heating
low expansion on
heating
low expansion on
heating
KMT of Matter
• According to the kinetic
molecular theory, the state of
a substance at room
temperature depends on the
strength of the attractions
between its particles.
Forces
• Intramolecular forces are
between the atoms within a
molecule = bonds
• Intermolecular forces are
between molecules
Intramolecular Forces
• Covalent Bonds – between
nonmetals; sharing of
electrons
• Ionic Bonds – between metals
and nonmetals; transfer of
electrons to form ions
• Metallic Bonds – between
metals; sea of electrons
Intermolecular Forces
1. Dipole-dipole attractions
a. Hydrogen bonding
2. Ion-dipole attractions
3. London dispersion forces
Dipole-Dipole Attractions
• Attractions due to permanent
dipoles in polar molecules
• Remember, a dipole is created
when positive and negative
charges are separated by
some distance.
• Only 1% as strong as covalent
or ionic bonds and even weaker
if distance between charges
increases.
Hydrogen Bonds
• Unusually strong dipole-dipole
attractions that occur among
molecules in which
hydrogen is bonded to
a highly
electronegative atom,
such as F, N, or O
Ion-Dipole Attractions
• Polar molecules surround ions
based on their
attraction for the
charge on the ion
London Dispersion Forces
• Explain the attraction that
exists between non-polar
molecules
• Even in these molecules the
electrons are not uniformly
distributed at every second
• Temporary dipolar
arrangement of charge
creates an instantaneous
dipole.
London Dispersion Forces
• Instantaneous dipoles can
induce similar dipoles in
neighboring atoms
Kinetic Energy
low
medium
high
Intermolecular
Attraction
high
medium
Kinetic
low
State at Room
Temp.
solids
liquids
gases
Kinetic energy determines if particles will
overcome the intermolecular forces keeping them
together. Thus, higher attractions mean higher
boiling points, because a higher kinetic energy will
be needed to overcome the attraction.
Properties of Solids
• Intermolecular forces keep
the particles of a solid packed
together tightly
- Particles are highly ordered
with fixed positions
- Only particle movement =
vibrations
• Liquids have similar distance
between particles
Bonding in Solids
2 basic types of solids based on
the nature of the particles
that make them up:
1. Crystalline
a. Atomic
b. Ionic
c. Covalent-network
2. Amorphous
Crystalline Solids
• Solid in which the
representative particles are in
a highly ordered, repeating
pattern called a crystal
• Can be studied as unit cells,
small representative units
that repeat throughout the
structure
Bonding in Crystalline
Solids
The physical properties of
solids, such as hardness,
electrical conductibility, and
melting point, depend on the
kind of particles that make up
the solid and on the strength
of the attractive forces
between them. (Includes ionic,
covalent, and atomic
substances)
Covalent-Network Solids
• A type of crystalline solid in
which strong covalent bonds
forma a network extending
throughout the solid
• Have very high melting points
due to strong covalent bonds
Amorphous Solids
• Solids in which the
arrangement of the
representative particles lacks
a regular, repeating pattern
• Also known as “supercooled
liquids”
- Liquids cooled until viscosity
becomes so high that no
flow can occur
Amorphous Solids
• Get softer when heated
instead of reaching an abrupt
melting point like crystalline
solids
• Examples: glass, rubber, some
plastics
Liquids
The physical properties of
liquids are determined by the
nature and strength of the
intermolecular forces present
between their molecules
Properties of Liquids
• Viscosity – resistance to motion that
exists between molecules of a liquid when
they move past each other
– Increased intermolecular forces yield
increased viscosity
• Glycerine has lots of
hydrogen bonds
• Decreased
temperatures yield
increased viscosity
Properties of Liquids
• Surface tension – imbalance of
attractive forces at the surface of a
liquid that cause the surface to behave
as if it had a film across it
»Explains “beading” of liquids
»As with viscosity, increased
intermolecular forces or
decreased
temperatures yield increased
surface tension
Phase Changes
• Changes of state are physical
changes, not chemical.
– Intermolecular forces break or
form, not intramolecular
• Phase changes occur when
energy enters or leaves a
compound.
• Energy in: solid g liquid g gas
• Energy out: gas g liquid g
solid
Heating/Cooling Curves
• Show the changes in
temperature as energy as heat
is added to (or removed from)
a substance and it changes
states.
Heating/Cooling Curve
Phase Changes
• During phase changes, heat is
still transferred, but no
temperature change takes
place.
Phase Diagrams
• Shows the state of matter of
a substance at increasing
pressures and temperatures.
• Can be used to determine the
state of matter of a
substance at a particular
pressure and temperature.
Phase Diagram of Water
Energy Requirements
• Molar heat of fusion: the energy required
to melt one mole of a solid
• Molar heat of vaporization: the energy
required to vaporize one mole of a liquid
• Molar heats of vaporization are higher
because more intermolecular forces will
have to be overcome to go from a liquid to
a gas than a solid to a liquid.
Vaporization
• Not all of the particles in a liquid have the
same kinetic energy.
• Temperature is a measurement of the
average KE.
• The particles of a liquid must be moving at
a sufficient speed to overcome the
intermolecular forces of the liquid to
escape as a gas.
• Thus, evaporation is endothermic: it
requires energy to occur.
Vapor Pressure
• Vapor pressure is the
pressure exerted by a vapor in
equilibrium with its liquid
phase at a certain
temperature.
• For a liquid in a closed
container, vaporization and
condensation will occur at an
equal rate once the equilibrium
vapor pressure is reached.
Vapor Pressure and Boiling Point
• When water begins to boil, bubbles of H2O
gas are created as some of the molecules
escape the intermolecular bonds holding
the liquid together.
• The bubbles will expand and rise to the
surface only if the pressure exerted by
the water vapor in the bubble is greater
than the atmospheric pressure.
Vapor Pressure and Boiling Point
• Thus, the vapor pressure of the water
must be equal to atmospheric pressure
before boiling can occur.
• At temperatures below 100oC, the vapor
pressure of water is below 1 atm, so the
atmospheric pressure prevents boiling from
occurring.
Elevation and Boiling Point
• What happens to atmospheric
pressure as you increase your
elevation above sea level?
• What would you expect to
happen to the boiling point of
water at high elevations?