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Chapter 11
Liquids and solidsLiquids and solids
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Almost all substances that are liquids are molecular, (held together by the covalent bonds within the molecule)
The physical properties of molecular liquids and solids is due to the intermolecular forces that hold them together
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They are similar compared to gases. They are incompressible. Their density doesn’t change with
temperature. These similarities are due
• to the molecules being close together in solids and liquids
• and far apart in gases What holds them close together?
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Intermolecular forces Inside molecules (intramolecular) the
atoms are bonded to each other. Intermolecular refers to the forces
between the molecules. These are what hold the molecules
together in the condensed states.
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Intermolecular forces are much weaker than ionic or covalent bonds
Less energy is required to vaporize water, than to break the bonds between hydrogen and oxygen.
When a molecule changes state from solid-liquid-gas, the molecule itself stays together
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Properties of liquids Many are related to the
intermolecular forces that hold them together
Boiling point- in order for a liquid to boil the molecules must overcome their attractive forces between them in order to separate. The stronger the intermolecular force, the higher the boiling point. (Melting point is based on the same relationship)
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Ion dipole forces Refer to a solution, not a molecule. Force that exists between an ion and
the partial charge on the end of a polar molecules.
Like NaCl dissolved in H2O The strength of the force increases
as the charge on the ion increases, or the strength of the dipole increases.
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Types of Intermolecular Forces3 Types for neutral molecules:
Dipole-dipole
London Dispersion
Hydrogen
All are based on electrostatic attractions
All called Van der Waals forces (man who predicted the deviation of gases from the ideal gas laws due to attractive forces between them)
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Dipole - Dipole 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.
1% as strong as covalent bonds Only work when polar molecules are
very close together
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For molecules of approximately equal mass and size, the strengths of intermolecular attractions increase with increasing polarity
The higher the dipole moment, the higher the boiling point
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London Dispersion Forces Non - polar molecules also exert
forces on each other. Otherwise, no solids or liquids. Electrons are not evenly distributed at
every instant in time. Have an instantaneous dipole. Induces a dipole in the atom next to it. Induced dipole- induced dipole
interaction.
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London Dispersion Forces Weak, short lived. Lasts longer at low temperature. Eventually long enough to make liquids. More electrons, more polarizable. Bigger molecules, higher melting and
boiling points. Much, much weaker than other forces. BP & MP of halogens and noble gases
increase as you go down
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Hydrogen Bonding 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 to get close. Effects boiling point.
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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
+-
+
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Example
H H H HH H H H
- +
H H H H
- +
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Liquids Many of the properties due to
internal attraction of atoms.
• Beading
• Surface tension
• Capillary action Stronger intermolecular forces cause
each of these to increase.
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Viscosity Which will pour very easily, syrup or
lemonade? Both liquids so what accounts for the difference?
Viscosity- How much a liquid resists flowing.
The higher the viscosity, the slower it flows.
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Related to how the molecules of the liquid move with respect to one another.
Viscosity decreases as temperature increases.
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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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Water has a high surface tension because of it’s hydrogen bonds
Cohesive- the intermolecular forces that bind like molecules together (very strong for water)
Adhesive- the intermolecular forces that bind a molecule to another surface.
Mercury/ water
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Capillary Action Liquids spontaneously rise in a
narrow tube. Helps water move through plants Occurs because of the attraction
between the polar glass and the polar 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.
When you wax
car.
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Relation to intermolecular forces? The stronger the force, the bigger
the effect.
• Hydrogen bonding
• Polar bonding
• LDF
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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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Phase changes Directly related to the strength of the
intermolecular forces> Na has a boiling point of 883°C
(strong metallic bonds) While iodine readily sublimes at
room temperature (LDF) When your phase change is going to
a less ordered state (more entropy) energy must be supplied
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Energy Requirement Fusion- to melt, the energy required
to change from the solid to liquid state is called the heat of fusion (Hfus)
Vaporization- to turn from a liquid to a gas, the energy required to change from a liquid to a gas is called the heat of vaporization (Hvap).
The heat of vaporization is usually much larger than the heat of fusion.
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Fusion, vaporization, and sublimation are endothermic
Freezing, condensation, and deposition are exothermic
Steam burns, sweating cools
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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.
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-40
-20
0
20
40
60
80
100
120
140
0 40 120 220 760 800
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 40 120 220 760 800
Heating Curve for Water
Heat of Fusion
Heat ofVaporization
Slope is Heat Capacity
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Calculating Energy Changes Q=mcΔT C= specific heat (given in J/g-K) M=mass ΔT= change in T, no need to convert
To melt To vaporize molesHfus molesHvap
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Calculate the enthalpy change of converting 1 mole of ice at -25°C to water vapor at 125°C. (The pressure is held constant at 1 atm). The specific heats of ice is 2.09 J/g-K, the specific heat of water is 4.18 J/g-K, and the specific heat of vapor is 1.84 J/g-K. The heat of fusion for water is 6.01 kJ/mol, and the heat of vaporization is 40.67 kJ/mol.
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You Try! Ethanol (C2H5OH) melts at -114°C,
and boils at 78°C. The enthalpy of fusion of ethanol is 5.02 kJ/mol, and the enthalpy of vaporization 38.56 kJ/mol. The specific heats of solid and liquid ethanol are respectively .97J/g-K, and 2.3J/g-K. How much heat is required to convert 75 g of ethanol at -120°C to the vapor phase at 78°C?
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Critical Temperature and Pressure As temperature rises, gases become
harder to liquefy because their molecules have high KE and are very far apart
Critical Temperature- the highest temperature a substance can exist as a liquid
Critical Pressure- The pressure required to liquefy the gas at the critical temperature.
The stronger the intermolecular forces, the easier a gas will liquefy.
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Vapor Pressure Vaporization - change from
liquid to gas at boiling point. Evaporation - change from
liquid to gas below boiling point
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Vaporization is an endothermic process - it requires heat.
Energy is required to overcome intermolecular forces pushing the molecules far apart into the gas phase.
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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.
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Dynamic Equilibrium The two opposing processes of
condensation and vaporization are occurring at the same rate.
May look as if nothing is happening because there is no net change
The vapor pressure is the pressure exerted by its vapor when the liquid and vapor states of a substance are at dynamic equilibrium
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In an open system As the water evaporates, the
molecules spread out and are not recaptured, no equilibrium is reached and the water continues to evaporate until there is nothing left.
Liquids that have high vapor pressure evaporate quickly and are called volatile (gasoline)
Vapor pressure increases with increasing T.
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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. Where will water boil faster, at the
top of a hill or the bottom of a valley? Where will pasta cook faster, at the
top of a hill or the bottom of a valley?
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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.
Shows where equilibrium exists between different states of matter.
Can use it to predict the most stable phase of matter at a given T and P.
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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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Solids Two major types. Amorphous- those with much
disorder in their structure. (Rubber and glass)
Crystalline- have a regular arrangement of components in their structure. (Diamond and Quartz)
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Melting Points The melting points of crystals are
definite. The melting points of amorphous
solids vary.
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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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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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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 there bonding effect their
properties.
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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 level 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- coal 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 sp3hybridized, 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?
Empty MOs
Filled MOs
EThe space between orbitals make it impossible for electrons to 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 through out these delocalized orbitals.
The layers slide by each other.
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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 size of atom. Large molecules (such as I2 ) can be
solids even without dipoles.
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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
4 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.
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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.
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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. As time goes by the rate of
vaporization remains constant but the rate of condensation increases because there are more molecules to condense.
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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. 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
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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
lnP
P =
H
R T-T
T1vap
T2vap
vap
2 1
1 1
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R = 8.3145 J/K mol. Surprisingly this is the graph of a
straight line. (actually the proof is in the book)
Mathematical relationship
lnP
P =
H
R T-T
T1vap
T2vap
vap
2 1
1 1