liquids and solids
DESCRIPTION
Chapter 10. Liquids and solids. They are similar to each other. Different than gases. They are incompressible. Their density doesn’t change much with temperature. These similarities are due to the molecules staying close together in solids and liquids, and far apart in gases. - PowerPoint PPT PresentationTRANSCRIPT
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Liquids and solidsLiquids and solids
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They are similar to each other Different than gases. They are incompressible. Their density doesn’t change much
with temperature. These similarities are due to the
molecules • staying close together in solids and
liquids, and • far apart in gases.
What holds them close together?
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Intermolecular forces (IMF’s) Inside molecules (intramolecular) the
atoms are bonded to each other. Intermolecular refers to the forces
between the molecules. Holds the molecules together in the
condensed states.
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Types of IMF’s Strong
• covalent bonding• ionic bonding
Weak• Dipole-dipole• London dispersion forces
During phase changes the molecules stay intact.
Energy is used to overcome these forces.
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Dipole - Dipole Remember where the polar definition
came from? 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.
Only 1% as strong as covalent bonds!! Even weaker with greater distance Play a small role in gases
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Hydrogen Bonding These are especially strong dipole-
dipole forces when H is attached to F, O, or N.
These three because-
• They have high electronegativity.
• They are small enough for the H to get close.
Affects boiling point. HOW????
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CH4
SiH4
GeH4SnH4
PH3
NH3 SbH3
AsH3
H2O
H2SH2Se
H2Te
HF
HI
HBrHCl
Boiling Points
0ºC
100
-100
200
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Water
+-
+
Each water molecule can make up to four H-bonds
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London Dispersion Forces Non-polar molecules also exert forces
on each other; otherwise, they would not make solids or liquids.
Electrons are not evenly distributed at every instant in time.
This phenomenon: • creates an instantaneous dipole.• induces a dipole in the atom next to it.
This creates an induced dipole- induced dipole interaction.
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Example
H H H HH H H H
+ -
H H H H
+ - +
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London Dispersion Forces Weak, short lived (“instantaneous”). Lasts longer at low temperature. WHY?? Eventually last long enough to make
liquids. More electrons = more polarizable. Bigger molecules = higher melting and
boiling points. Weaker than other forces.
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Van der Waal’s forces London dispersion forces and Dipole
interactions Order of increasing strength
• LDF
• Dipole-dipole
• H-bond
• “Real” bonds
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Liquids Many of the properties are due to
internal attraction of atoms.
• Beading
• Surface tension
• Capillary action
• Viscosity Stronger intermolecular forces cause
each of these to increase.
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Surface tension
Molecules in the middle are attracted in all directions.
Molecules at the the top are only pulled inside.
Minimizes surface area.
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Capillary Action Liquids spontaneously rise in a
narrow tube. Intermolecular forces are cohesive,
connecting like things. Adhesive forces connect to
something else. Glass is polar.
• It attracts water molecules.
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Beading If a polar substance
is placed on a non-polar surface.
• There are cohesive,
• But no adhesive forces.
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Viscosity How much a liquid resists flowing. Large forces, more viscous. Large molecules can get tangled up. Cyclohexane has a lower viscosity
than hexane. Because it is a circle- more compact.
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How much of these? Stronger forces, bigger effect.
• Hydrogen bonding
• Dipole-dipole
• LDF In that order
•H next to O,N, or F
•Polar molecules
•All molecules
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Model of a Liquid Can’t see molecules so picture them
as:
1. In motion but attracted to each other
2. With regions arranged like solids but:
• with higher disorder.
• with fewer holes than a gas.
• highly dynamic - regions changing between types.
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Phases The phase of a substance is
determined by three things.
• The temperature.
• The pressure.
• The strength of intermolecular forces.
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Solids Two major types. Amorphous- those with much
disorder in their structure. Crystalline- have a regular
arrangement of components in their structure.
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Crystals Lattice- a three dimensional grid that
describes the locations of the pieces in a crystalline solid.
Unit Cell-The smallest repeating unit in of the lattice.
Three common types.
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Cubic
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Body-Centered Cubic
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Face-Centered Cubic
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The book drones on about Using diffraction patterns to identify
crystal structures. Talks about metals and the closest
packing model. It is interesting, but trivial. We need to focus on metallic bonding. Why do metal atoms stay together? How their bonding affects their
properties.
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Solids There are many amorphous solids. Like glass. We tend to focus on crystalline solids. two types.
• Ionic solids have ions at the lattice points.
• Molecular solids have molecules. Sugar vs. Salt.
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Metallic Bonds How atoms are held together in the
solid. Metals hold onto their valence
electrons very weakly. Think of them as positive ions
floating in a sea of electrons.
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Sea of Electrons
+ + + ++ + + +
+ + + +
Electrons are free to move through the solid.
Metals conduct electricity.
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Metals are Malleable Hammered into shape (bend). Ductile - drawn into wires. Because of mobile valence electrons
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Malleable
+ + + ++ + + +
+ + + +
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Malleable
+ + + +
+ + + ++ + + +
Electrons allow atoms to slide by but still be attracted.
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Metallic bonding
1s
2s2p
3s
3pFilled Molecular Orbitals
Empty Molecular Orbitals
Magnesium Atoms
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Filled Molecular OrbitalsEmpty Molecular Orbitals
The 1s, 2s, and 2p electrons are close to nucleus, so they are not able to move around.
1s
2s2p
3s
3p
Magnesium Atoms
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Filled Molecular OrbitalsEmpty Molecular Orbitals
1s
2s2p
3s
3p
Magnesium Atoms
The 3s and 3p orbitals overlap and form molecular orbitals.
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Filled Molecular OrbitalsEmpty Molecular Orbitals
1s
2s2p
3s
3p
Magnesium Atoms
Electrons in these energy levels can travel freely throughout the crystal.
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Filled Molecular OrbitalsEmpty Molecular Orbitals
1s
2s2p
3s
3p
Magnesium Atoms
This makes metals conductors
Malleable because the bonds are flexible.
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Carbon- A Special Atomic Solid There are three types of solid carbon. Amorphous- soot - uninteresting. Diamond- hardest natural substance
on earth, insulates both heat and electricity.
Graphite- slippery, conducts electricity.
How the atoms in these network solids are connected explains why.
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Diamond- each Carbon is sp3
hybridized, connected to four other carbons.
Carbon atoms are locked into tetrahedral shape.
Strong bonds give the huge molecule its hardness.
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Why is it an insulator?
All the electrons need to be shared in the covalent bonds
Can’t move around
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Each carbon is connected to three other carbons and sp2 hybridized.
The molecule is flat with 120º angles in fused 6 member rings.
The bonds extend above and below the plane.
Graphite is different.
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This bond overlap forms a huge bonding network.
Electrons are free to move throughout these delocalized orbitals.
Conducts electricity
The layers slide by each other.
Lubricant
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Molecular solids. Molecules occupy the corners of the
lattices. Different molecules have different
forces between them. These forces depend on the size of
the molecule. They also depend on the strength
and nature of dipole moments.
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Those without dipoles.
Most are gases at 25ºC. The only forces are London Dispersion
Forces. These depend on number of electrons. Large molecules (such as I2 ) can be
solids even without dipoles. (LDF)
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Those with dipoles. Dipole-dipole forces are generally
stronger than L.D.F. Hydrogen bonding is stronger than
Dipole-dipole forces. No matter how strong the
intermolecular force, it is always much, much weaker than the forces in bonds.
Stronger forces lead to higher melting and freezing points.
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Water is special
HO
H
-
Each molecule has two polar O-H bonds.
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Water is special
HO
H
Each molecule has two polar O-H bonds.
Each molecule has two lone pair on its oxygen.
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Water is special Each molecule has two polar
O-H bonds. Each molecule has two lone
pair on its oxygen. Each oxygen can interact with
2 hydrogen atoms.
HO
H
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Water is special
HO
H
HO
H
HO
H
This gives water an especially high melting and boiling point.
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Ionic Solids The extremes in dipole-dipole forces-
atoms are actually held together by opposite charges.
Huge melting and boiling points. Atoms are locked in lattice so hard and
brittle. Every electron is accounted for so they
are poor conductors-good insulators. Until melted or dissolved.
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Phase Changes
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Vapor Pressure Vaporization - change from
liquid to gas at boiling point. Evaporation - change from
liquid to gas below boiling point
Heat (or Enthalpy) of
Vaporization (Hvap )- the
energy required to vaporize
1 mol at 1 atm.
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Vaporization is an endothermic process - it requires heat.
Energy is required to overcome intermolecular forces.
Responsible for cool beaches. Why we sweat.
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Condensation Change from gas to liquid. Achieves a dynamic equilibrium with
vaporization in a closed system. What is a closed system? A closed system means matter
can’t go in or out. Put a cork in it. What the heck is a “dynamic
equilibrium?”
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When first sealed the molecules gradually escape the surface of the liquid
Dynamic equilibrium
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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
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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
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Rate of Vaporization = Rate of
Condensation Molecules are constantly changing
phase “Dynamic” The total amount of liquid and vapor
remains constant “Equilibrium”
Dynamic equilibrium
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Vapor pressure The pressure above the liquid at
equilibrium. Liquids with high vapor pressures
evaporate easily. They are called volatile. Decreases with increasing
intermolecular forces.
• Bigger molecules (bigger LDF)
• More polar molecules (dipole-dipole)
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Vapor pressure Increases with increasing
temperature. Easily measured in a barometer.
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Dish of Hg
Vacuum
Patm=
760 torr
A barometer will hold a column of mercury 760 mm high at one atm
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Dish of Hg
Vacuum
Patm=
760 torr
A barometer will hold a column of mercury 760 mm high at one atm.
If we inject a volatile liquid in the barometer it will rise to the top of the mercury.
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Dish of Hg
Patm=
760 torr
A barometer will hold a column of mercury 760 mm high at one atm.
If we inject a volatile liquid in the barometer it will rise to the top of the mercury.
There it will vaporize and push the column of mercury down.
Water
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Dish of Hg
736 mm Hg
Water Vapor
The mercury is pushed down by the vapor pressure.
Patm = PHg + Pvap
Patm - PHg = Pvap
760 - 736 = 24 torr
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Temperature Effect
Kinetic energy
# of
mol
ecu
les
T1
Energy needed to overcome intermolecular forces
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Kinetic energy
# of
mol
ecu
les
T1
Energy needed to overcome intermolecular forces
T1
T2
At higher temperature more molecules have enough energy - higher vapor pressure.
Energy needed to overcome intermolecular forces
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Mathematical relationship
ln is the natural logarithm
• ln = Log base e
• e = Euler’s number an irrational number like
Hvap is the heat of vaporization in J/mol
1
2
vapT vap
vapT 2 1
P H 1 1ln = -
P R T T
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R = 8.3145 J/K mol. Surprisingly this is the graph of a
straight line. If you graph ln P vs 1/T
Mathematical relationship1
2
vapT vap
vapT 2 1
P H 1 1ln = -
P R T T
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The vapor pressure of water is 23.8 torr at 25°C. The heat of vaporization of water is 43.9 kJ/mol. Calculate the vapor pressure at 50°C
At what temperature would it have a vapor pressure of 760 torr?
Mathematical relationship1
2
vapT vap
vapT 2 1
P H 1 1ln = -
P R T T
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Changes of state The graph of temperature versus
heat applied is called a heating curve.
The temperature a solid turns to a liquid is the melting point.
The energy required to accomplish this change is called the Heat (or Enthalpy) of Fusion Hfus
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-40
-20
0
20
40
60
80
100
120
140
0 10 90 190 730 740
Heating Curve for Water
IceWater and Ice
Water
Water and Steam
Steam
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-40
-20
0
20
40
60
80
100
120
140
0 10 90 190 730 740
Heating Curve for Water
Heat of Fusion
Heat ofVaporization
1/ Slope is Heat Capacity
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Melting Point Melting point is determined by the
vapor pressure of the solid and the liquid.
At the melting point the vapor pressure of the solid =
vapor pressure of the liquid
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Solid Water
Liquid Water
Water Vapor Vapor
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Solid Water
Liquid Water
Water Vapor Vapor
If the vapor pressure of the solid is higher than that of the liquid the solid will release molecules to achieve equilibrium.
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Solid Water
Liquid Water
Water Vapor Vapor
While the molecules of condense to a liquid.
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This can only happen if the temperature is above the freezing point since solid is turning to liquid.
Solid Water
Liquid Water
Water Vapor Vapor
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If the vapor pressure of the liquid is higher than that of the solid, the liquid will release molecules to achieve equilibrium.
Solid Water
Liquid Water
Water Vapor Vapor
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Solid Water
Liquid Water
Water Vapor Vapor
While the molecules condense to a solid.
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The temperature must be below the freezing point since the liquid is turning to a solid.
Solid Water
Liquid Water
Water Vapor Vapor
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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.
Solid Water
Liquid Water
Water Vapor Vapor
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Boiling Point Reached when the vapor pressure
equals the external pressure. Normal boiling point is the boiling
point at 1 atm pressure. Superheating - Heating above the
boiling point. Supercooling - Cooling below the
freezing point.
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Phase Diagrams. A plot of temperature versus
pressure for a closed system, with lines to indicate where there is a phase change.
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Temperature
SolidLiquid
Gas
1 Atm
AA
BB
CCD
D D
Pre
ssur
e
D
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SolidLiquid
Gas
Triple Point
Critical Point
Temperature
Pre
ssur
e
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SolidLiquid
Gas
This is the phase diagram for water. The density of liquid water is higher
than solid water.
Temperature
Pre
ssur
e
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Solid Liquid
Gas
1 Atm
This is the phase diagram for CO2
The solid is more dense than the liquid The solid sublimes at 1 atm.
Temperature
Pre
ssur
e
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