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Liquids and solids
Comparing Solids, Liquids,
Gases according to the Kinetic
Theory
They are similar to each other
u Different
than gases.
are incompressible.
u Their density doesn’t change much
with temperature.
u These similarities are due
• to the molecules staying close
together in solids and liquids
• and far apart in gases
u What holds them close together?
u They
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Phase Transitions
Increasing Thermal Energy (Temperature))
Sublimation
Fusion
(Melting)
Evaporation
Crystallization
(Freezing)
Condensation
Deposition
Decreasing Thermal Energy (Temperature))
Vapor Pressure
Vaporization - change from
liquid to gas at boiling point.
u Evaporation - change from
liquid to gas below boiling
point
u Heat (or Enthalpy) of
u
Vaporization (DHvap )- the
energy required to vaporize
1 mol at 1 atm.
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u Vaporization
is an endothermic
process - it requires heat.
u Energy is required to overcome
intermolecular forces.
u Responsible for cool beaches.
u Why we sweat.
Condensation
u Change
from gas to liquid.
a dynamic equilibrium with
vaporization in a closed system.
u What is a closed system?
u A closed system means matter
can’t go in or out.
u Put a cork in it.
u What the heck is a “dynamic
equilibrium?”
u Achieves
Dynamic equilibrium
/When
first sealed the molecules
gradually escape the surface of
the liquid
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Dynamic equilibrium
/When
first sealed the
molecules gradually escape
the surface of the liquid
/As the molecules build up
above the liquid some
condense back to a liquid.
Dynamic equilibrium
/As
time goes by the rate of
vaporization remains constant
/ but the rate of condensation
increases because there
are more molecules to
condense.
/Equilibrium is reached
when
Dynamic equilibrium
Rate of Vaporization =
Rate of Condensation
u Molecules
are constantly changing
phase “Dynamic”
u The total amount of liquid and vapor
remains constant “Equilibrium”
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Vapor pressure
u The
pressure above the liquid at
equilibrium.
u Liquids with high vapor pressures
evaporate easily.
u They are called volatile.
u Decreases with increasing
intermolecular forces.
• Bigger molecules (bigger LDF)
• More polar molecules (dipole-dipole)
Vapor pressure
u Increases
with increasing
temperature.
u Easily measured in a barometer.
Changes of state
u The
graph of temperature versus
heat applied is called a heating
curve.
u The temperature a solid turns to a
liquid is the melting point.
u The energy required to accomplish
this change is called the Heat (or
Enthalpy) of _______
DH___
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Energy Associated with Changes of State
Change of State
Melting
Freezing
Evaporation
Condensation
Name of Energy
DHfus
DHcrys
DHvap
DHcond
Heat
Heat
Heat
Heat
of
of
of
of
Fusion
Crystallization
Vaporization
Condensation
DHXXX = mC
where C = Molar Heat of ___ Constant J/mol
m = moles of substance
Heating Curve for Water
Water and
Steam
Steam
Water
Ice
Water and
Ice
Calculating Heat of Phase Change
DHXXX = mC
C = Molar Heat of ___ Constant J/mol
m = moles of substance
For water:
Molar Heat of Fusion Constant C = 6.01 kJ/mol
Molar Heat of Vaporization Constant C = 40.7 kJ/mol
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Calculating Heat of Phase Change
DHXXX = mC
C = Molar Heat of ___ Constant J/mol
m = moles of substance
For water:
Molar Heat of Fusion Constant C = 6.01 kJ/mol
Molar Heat of Vaporization Constant C = 40.7 kJ/mol
Calculate the amount of heat needed to melt 35.0 g of ice at 0 ºC.
Express your answer in kilojoules.
Calculating Heat of Phase Change
DHXXX = mC
C = Molar Heat of ___ Constant J/mol
m = moles of substance
For water:
Molar Heat of Fusion Constant C = 6.01 kJ/mol
Molar Heat of Vaporization Constant C = 40.7 kJ/mol
How much energy is released to the environment by 50.0 grams of
condensing water vapor?
How much energy is needed to change
20.0g of ice at -10.0 ⁰C to water at 60.0⁰C?
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Melting Point
u Melting
point is determined by the
vapor pressure of the solid and the
liquid.
u At the melting point the
vapor pressure of the solid = vapor pressure of the liquid
Water Vapor
Vapor
Solid
Water
Liquid
Water
u If
the vapor pressure of the solid is higher
than that of the liquid the solid will
release molecules to achieve equilibrium.
Water Vapor
Vapor
Solid
Water
Liquid
Water
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u While
the molecules of condense
to a liquid.
Water Vapor
Vapor
Solid
Water
Liquid
Water
u This
can only happen if the temperature
is above the freezing point since solid
is turning to liquid.
Water Vapor
Vapor
Solid
Water
Liquid
Water
u If
the vapor pressure of the liquid is higher
than that of the solid, the liquid will
release molecules to achieve equilibrium.
Water Vapor
Vapor
Solid
Water
Liquid
Water
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u While
the molecules condense to
a solid.
Water Vapor
Vapor
Solid
Water
Liquid
Water
u The
temperature must be below
the freezing point since the liquid
is turning to a solid.
Water Vapor
Vapor
Solid
Water
Liquid
Water
u If
the vapor pressure of the solid and
liquid are equal, the solid and liquid are
vaporizing and condensing at the same
rate. The Melting point.
Water Vapor
Vapor
Solid
Water
Liquid
Water
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Boiling Point
u Reached
when the vapor pressure
equals the external pressure.
u Normal boiling point is the boiling
point at 1 atm pressure.
u Superheating - Heating above the
boiling point.
u Supercooling - Cooling below the
freezing point.
Intermolecular forces
u Inside
molecules (intramolecular) the
atoms are bonded to each other.
u Intermolecular refers to the forces
between the molecules.
u Holds the molecules together in the
condensed states…
• i.e. liquids and solids
Intermolecular forces-Liquids
• Hydrogen Bonding
• Dipole dipole
• London dispersion forces
• Van Der Waals
u During phase changes the molecules
stay intact.
u Energy used to overcome
intermolecular forces (IMF)
• The stronger the IMF, more energy
needed.
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Hydrogen Bonding-Liquids
u Especially
strong dipole-dipole forces
when H is attached to F, O, or N
u These three because• They have high electronegativity.
• They are small enough to get close.
u Effects boiling point.
Water
d+
dd+
Each water molecule can make up to four
H-bonds
Water is special
u Each
molecule has two polar
O-H bonds.
d- O
d+
H
H
d+
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Water is special
u Each
O
molecule has two polar
O-H bonds.
d+ u Each molecule has two lone
H pair on its oxygen.
H
d+
Water is special
u Each
molecule has two polar
O-H bonds.
d+ u Each molecule has two lone
H pair on its oxygen.
O
u Each oxygen can interact with
H 2 hydrogen atoms.
d+
Water is special
O
d+
H
H
d+
O
d+
H
O
d+
H
H
d+
u This
gives water
an especially
high melting
and boiling
point.
H
d+
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100 H2O
0ºC
Boiling Points
HF
NH3
H2S
HCl
-100
PH3
SiH4
200
H2Te
SbH3
HI
H2Se
AsH3
HBr
GeH4
SnH4
CH4
Dipole – Dipole Liquids
u Remember
where the polar definition
came from?
u Molecules line up in the presence of a
electric field. The opposite ends of
the dipole can attract each other so
the molecules stay close together.
u 1% as strong as covalent bonds
u Weaker with greater distance.
u Small role in gases.
-
+
+
+
-
-
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London Dispersion Forces Liquids
u Non
- polar molecules also exert
forces on each other.
u Otherwise, no solids or liquids.
u Electrons are not evenly distributed at
every instant in time.
u Have an instantaneous dipole.
u Induces a dipole in the atom next to it.
u Induced dipole- induced dipole
interaction.
London Dispersion Forces
Example
d+
d-
d+
d-
H
H
H
H
London Dispersion Forces
u Weak,
short lived.
u Lasts longer at low temperature.
u Eventually long enough to make liquids.
u More electrons, more polarizable.
u Bigger molecules, higher melting and
boiling points.
u Weaker than Hydrogen bonding and
Dipole dipole forces.
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Van der Waal’s forces Liquids
u Intermolecular
force between
nonpolar molecules
u Order of increasing strength
• Van der Waals
• LDF
• Dipole
• H-bond
• Intramolecular bonds
Liquids Properties
u Many
of the properties due to
internal attraction of molecules.
• Beading
• Surface tension
• Capillary action
• Viscosity
u Stronger intermolecular forces cause
each of these to increase.
Liquids…Surface tension
u Molecules
at
the the top are
only pulled
inside.
u Molecules in
the middle are
attracted in all
directions.
u Minimizes
surface area.
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Liquids…Capillary Action
u Liquids
spontaneously rise in a
narrow tube.
u Intermolecular forces are cohesive,
connecting like things.
u Adhesive forces connect to
something else.
u Glass is polar.
• It attracts water molecules.
Liquids…Beading
u If
a polar substance
is placed on a nonpolar surface.
• There are cohesive,
• But no adhesive
forces.
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Liquids…Viscosity
u How
much a liquid resists flowing.
u Large forces, more viscous.
u Large molecules can get tangled up.
u Cyclohexane has a lower viscosity
than hexane.
u Because it is a circle- more compact.
Liquids Properties
u Many
of the properties due to
internal attraction of molecules.
• Vapor Pressure
u The higher the molecular weight
causes the vapor pressure to
decrease.
Liquids Summary
u Many
of the properties due to
molecular weight when only London
Dispersion Forces are acting
• London Dispersion Force
• Boiling Point
• Surface Tension
• Viscosity
u The higher the molecular weight
causes each of these to increase.
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Liquids Summary
u Many
of the properties due to
molecular weight when only London
Dispersion Forces are acting
• Vapor Pressure
u The higher the molecular weight
causes the vapor pressure to
decrease.
Liquids Summary
u Stronger
forces, bigger effect.
• Hydrogen bonding •H next to O,N, or F
• Dipole-dipole
•Polar molecules
• LDF
•Nonpolar (only), all
• Van der Waals
•Nonpolar molecules
u In that order
Summary of Properties of
Liquids pg 434
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Solids
u Two major
types.
• Amorphous- those with much disorder
in their structure.
• Crystalline- have a regular
arrangement of components in their
structure.
Solids
u There
are many amorphous solids.
u Like glass.
u We tend to focus on crystalline solids.
u two types.
• Ionic solids have ions at the lattice
points.
• Molecular solids have molecules.
u Sugar vs. Salt.
Crystals
u Lattice-
a three dimensional grid that
describes the locations of the pieces
in a crystalline solid.
u Unit Cell-The smallest repeating unit
in of the lattice.
u Three common types.
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Cubic
Face-Centered Cubic
Body-Centered Cubic
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Types of Structures-Solids
• Molecular Solids
– Neon, ice, dry ice
• Metallic Solids
– Iron, copper, silver
• Ionic Solids
– Sodium chloride, zinc sulfide
• Covalent Network Solids
– Diamond, graphite
Molecular solids.
u Molecules
occupy the corners of the
lattices.
u Different molecules have different
forces between them.
u These forces depend on the size of
the molecule.
u They also depend on the strength
and nature of dipole moments.
Molecular Solids without dipoles.
Most are gases at 25ºC.
u The only forces are London Dispersion
Forces.
u These depend on number of electrons.
u Large molecules (such as I2 ) can be
solids even without dipoles. (LDF)
u
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Molecular Solids with dipoles.
u Dipole-dipole
forces are generally
stronger than L.D.F.
u Hydrogen bonding is stronger than
Dipole-dipole forces.
u No matter how strong the
intermolecular force, it is always
much, much weaker than the forces in
bonds.
u Stronger forces lead to higher melting
and freezing points.
Metallic Bonds
u How
atoms are held together in the
solid.
u Metals hold onto their valence
electrons very weakly.
u Think of them as positive ions
floating in a sea of electrons.
Metallic Bonds Sea of Electrons
u Electrons
are free to move through
the solid.
u Metals conduct electricity.
+
+ + +
+ + + +
+ + + +
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Ionic Solids
The extremes in dipole-dipole forces-atoms
are actually held together by opposite
charges.
u Huge melting and boiling points.
u Atoms are locked in lattice so hard and
brittle.
u Every electron is accounted for so they are
poor conductors-good insulators.
u Until melted or dissolved- goodconductor.
u
Covalent Network
Carbon- A Special Atomic Solid
There are three types of solid carbon.
Amorphous- soot - uninteresting.
u Diamond- hardest natural substance on
earth, insulates both heat and electricity.
u Graphite- slippery, conducts electricity.
u How the atoms in these network solids are
connected explains why.
u
u
Covalent Network
Diamondeach Carbon is sp3
hybridized, connected
to four other carbons.
Carbon atoms are
locked into
tetrahedral shape.
u Strong s bonds give
the huge molecule its
hardness.
u
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Covalent Network
Diamond-Why is it an insulator?
All the electrons
need to be shared
in the covalent
bonds
Can’t move around
Graphite is different.
Covalent Network
Each carbon is
connected to three
other
carbons and
sp2 hybridized.
u The molecule is flat
with 120º angles in
fused 6 member rings.
u The p bonds extend above and below the
plane.
u
This p bond overlap forms a huge p
bonding network. Covalent
Network
Electrons are free to move throughout
these delocalized orbitals.
u Conducts
electricity
u The layers slide
by each other.
u Lubricant
u
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Types of Structures-Solids
Effect on Melting Point
• Molecular Solids
– Weak, low melting point
• Ionic Solids and Covalent Network
Solids
– Stronger IMF, higher melting points
– Ionic Solids, greater the lattice energy, higher the
melting point
• Metallic Solids
– IA, IIA low melting point
– Increases melting point to VIIIB
– Decreases melting point again IB, IIB
Types of Structures-Solids
Effect on Hardness
• Molecular Solids
– Weak, soft and brittle
• Ionic Solids
– Strong, hard and brittle
• Covalent Network Solids
– Strongest, hardest
• Metallic Solids
– Not brittle, malleable
Metals are Malleable
degree of hardness
u Hammered
into shape (bend).
- drawn into wires.
u Because of mobile valence electrons
u Ductile
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Malleable
u Electrons
allow atoms to slide by but
still be attracted.
+ + + +
+ + + +
+ + + +
Malleable
+
+ + +
+ + + +
+ + + +
Types of Structures-Solids
Effect on Conductivity
• Molecular Solids
– nonconducting
• Metallic Solids
– Conducting (sea of electrons)
• Ionic Solids
– conducting only as liquids (ions)
• Covalent Network Solids
– nonconducting
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Summary of Solid Properties
Phase Diagrams
Phase Diagrams.
uA
plot of temperature versus
pressure for a closed system, with
lines to indicate where there is a
phase change.
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Pressure
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D
Solid
D
Liquid
1 Atm
B
C
A
C
D
B
D
A
Gas
Pressure
Temperature
Solid
Liquid
Critical
Point
Triple
Point
Gas
Temperature
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