Mr. Hollister HollidayLegacy High School
Regular & Honors Chemistry
Liquids & Solids
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Liquids
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Properties of the States of Matter:Liquids
• High densities compared to gases.• Fluid.
The material exhibits a smooth, continuous flow as it moves.
• Take the shape of their container(s).
• Keep their volume, do not expand to fill their container(s).
• Cannot be compressed into a smaller volume.
Particles in a Liquid
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Liquids• The particles in a liquid are closely
packed, but they have some ability to move around.
• The close packing results in liquids being incompressible.
• But the ability of the particles to move allows liquids to take the shape of their container and to flow. However, they don’t have enough freedom to escape and expand to fill the container(s).
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Properties of Liquids:Surface Tension
• Liquids tend to minimize their surface—a phenomenon we call surface tension.
• This tendency causes liquids to have a surface that resists penetration.
• The stronger the attractive force between the molecules, the larger the surface tension.
Water Strider
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Properties of Liquids:Viscosity
• Some liquids flow more easily than others.
• The resistance of a liquid’s flow is called viscosity.
• The stronger the attractive forces between the molecules, the more viscous the liquid is.
• Also, the less round the molecule’s shape, the larger the liquid’s viscosity.Some liquids are more viscous because
their molecules are long and get tangled in each other, causing them to resist flowing.
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Evaporation• Over time, liquids evaporate—the molecules of
the liquid mix with and dissolve in the air.• The evaporation happens at the surface.• Molecules on the surface experience a smaller
net attractive force than molecules in the interior.
• All the surface molecules do not escape at once, only the ones with sufficient kinetic energy to overcome the attractions will escape.
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Escaping from the Surface• The process of molecules of a
liquid breaking free from the surface is called evaporation.Also known as vaporization.
• Evaporation is a physical change in which a substance is converted from its liquid form to its gaseous form.The gaseous form is called a
vapor.
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Factors Effecting the Rate of Evaporation
• Liquids that evaporate quickly are called volatile liquids, while those that do not are called nonvolatile.
• Increasing the surface area increases the rate of evaporation.More surface molecules.
• Increasing the temperature increases the rate of evaporation.Raises the average kinetic energy,
resulting in more molecules that can escape.
• Weaker attractive forces between the molecules = faster rate of evaporation.
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Reconnecting with the Surface• When a liquid evaporates in a
closed container, the vapor molecules are trapped.
• The vapor molecules may eventually bump into and stick to the surface of the container or get recaptured by the liquid. This process is called condensation.A physical change in which a
gaseous form is converted to a liquid form.
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Dynamic Equilibrium• Evaporation and
condensation are opposite processes.
• Eventually, the rate of evaporation and rate of condensation in the container will be the same.
• Opposite processes that occur at the same rate in the same system are said to be in dynamic equilibrium.
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Evaporation and Condensation
Eventually, the condensation and
evaporation reach the same speed. The air in the flask is
now saturated with water vapor.
Shortly, the waterstarts to evaporate.
Initially the rateof evaporation is
much faster than rate of condensation
When water is justadded to the flask and it is capped, all the water molecules
are in the liquid.
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Vapor Pressure• Once equilibrium is reached,
from that time forward, the amount of vapor in the container will remain the same.As long as you don’t change
the conditions.• The partial pressure exerted
by the vapor is called the vapor pressure.
• The vapor pressure of a liquid depends on the temperature and strength of intermolecular attractions.
Vapor Pressure Curves
• The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure.
• The normal boiling point is the temperature at which its vapor pressure is 760 torr or 1 atm.
Water: A Unique and Important Substance
• Water is found in all three states on Earth.
• As a liquid, it is the most common solvent found in nature.
• Without water, life as we know it could not exist.The search for extraterrestrial
life starts with the search for water.
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Water
• Liquid at room temperature.Most molecular substances that have a molar
mass (18.02 g/mol) similar to water’s are gaseous.
• Relatively high boiling point.• Expands as it freezes.
Most substances contract as they freeze.Causes ice to be less dense than liquid water.
Solids
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Properties of the States of Matter:Solids
• High densities compared to gases.• Nonfluid.
They move as an entire “block” rather than a smooth, continuous flow.
• Keep their own shape, do not take the shape of their container(s).
• Keep their own volume, do not expand to fill their container(s).
• Cannot be compressed into a smaller volume.
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Solids• Some solids have their particles
arranged in an orderly geometric pattern. We call these crystalline solids.Salt and diamonds.
• Other solids have particles that do not show a regular geometric pattern over a long range. We call these amorphous solids.Plastic and glass.
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Crystalline Solids
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Types of Crystalline Solids
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Molecular Crystalline Solids
• Molecular solids are solids whose composite units are molecules.
• Solid held together by intermolecular attractive forces.Dispersion, dipole-dipole,
or H-bonding.• Generally low melting
points
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Ionic Crystalline Solids• Ionic solids are solids whose
composite units are formula units.• Solid held together by
electrostatic attractive forces between cations and anions.Cations and anions arranged in a
geometric pattern called a crystal lattice to maximize attractions.
• Generally higher melting points than molecular solids.Because ionic bonds are stronger
than intermolecular forces.
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Atomic Crystalline Solids
• Atomic solids are solids whose composite units are individual atoms.
• Solids held together by either covalent bonds, dispersion forces, or metallic bonds.
• Melting points vary depending on the attractive forces between the atoms.
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Practice—Classify Each of the Following Crystalline Solids as Molecular, Ionic, or Atomic.
• H2O(s)
• Si(s)
• C12H22O11(s)
• CaF2(s)
• Sc(NO3)3(s)
• H2O(s)—molecular.
• Si(s)—atomic.
• C12H22O11(s)—molecular.
• CaF2(s)—ionic.
• Sc(NO3)3(s)—ionic.
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Types of Atomic Solids
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Types of Atomic Solids:Covalent Network
• Covalent atomic solids have their atoms attached by covalent bonds. Effectively, the entire solid is one giant molecule.
• Because covalent bonds are strong, these solids have very high melting points.
• Because covalent bonds are directional, these substances tend to be very hard.
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Types of Atomic Solids:Nonbonding
• Nonbonding atomic solids are held together by dispersion forces.
• Because dispersion forces are relatively weak, these solids have very low melting points
• All the noble gases form nonbonding atomic solids.
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Types of Atomic Solids:Metallic
• Metallic solids are held together by metallic bonds.
• Metal atoms release some of their electrons to be shared by all the other atoms in the crystal.
• The metallic bond is the attraction of the metal cations for the mobile electrons.Often described as islands of
cations in a sea of electrons.
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Metallic Bonding• The model of metallic bonding
can be used to explain the properties of metals.
• The luster, malleability, ductility, and electrical and thermal conductivity are all related to the mobility of the electrons in the solid.
• The strength of the metallic bond varies, depending on the charge and size of the cations, so the melting points of metals vary as well.
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Practice—Decide if Each of the Following Atomic Solids Is Covalent, Metallic, or
Nonbonding.
• diamond
• neon
• iron
• diamond covalent.
• neon nonbonding.
• iron metallic.
Phase Changes
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Heating Curve Diagrams
• The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other.
• The temperature of the substance does not rise during the phase change.
Phase DiagramsPhase diagrams display the state of a
substance at various pressures and temperatures.
Phase Diagrams
• The AB line is the liquid-vapor interface.• It starts at the triple point (A), the point at
which all three states are in equilibrium.
Phase DiagramsIt ends at the critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.
Phase DiagramsEach point along this line is the boiling point of the substance at that pressure.
Phase Diagrams• The AD line is the interface between liquid and
solid.• The melting point at each pressure can be
found along this line.
Phase Diagrams• Below A the substance cannot exist in the
liquid state.• Along the AC line the solid and gas phases are
in equilibrium; the sublimation point at each pressure is along this line.
Phase Diagram of CO2
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Sublimation
• Sublimation is a physical change in which the solid form changes directly to the gaseous form.Without going through the
liquid form.
• Like melting, sublimation is endothermic.
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Intermolecular Forces
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Why Are Molecules Attracted to Each Other?
• Intermolecular attractions are a result of attractive forces between opposite charges.
• + ion to – ion.• + end of one polar molecule to − end of another
polar molecule.H-bonding is especially strong.Even nonpolar molecules will have temporary induced
dipoles.• Larger charge = stronger attraction.
Intermolecular Forces
The attractions between molecules are not nearly as strong as the intramolecular attractions that
hold compounds together.
Intermolecular Forces
They are, however, strong enough to control physical properties such as boiling and melting
points, vapor pressures, and viscosities.
Intermolecular Forces
These intermolecular forces as a group are referred to as van der Waals forces.
van der Waals Forces• London dispersion forces• Dipole-dipole interactions• Hydrogen bonding
Johannes van der Waals
Types of Intermolecular ForcesType of
force
Relative
strength
Present
in Example
London
Dispersion
force
Weak, but
increases
with molar
mass
All atoms
and
molecules
H2
Dipole–
Dipole
force
Moderate
Only
polar
molecules
HCl
Hydrogen
BondStrong
Molecules
having H
bonded to
F, O, or N
HF
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Dispersion Forces• Also known as London forces or
instantaneous dipoles.• Caused by distortions in the electron cloud of
one molecule inducing distortion in the electron cloud on another.
• Distortions in the electron cloud lead to a temporary dipole.
• The temporary dipoles lead to attractions between molecules—dispersion forces.
• All molecules have attractions caused by dispersion forces.
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Instantaneous Dipoles
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Strength of the Dispersion Force
• Depends on how easily the electrons can move, or be polarized.
• The more electrons and the farther they are from the nuclei, the larger the dipole that can be induced.
• Strength of the dispersion force gets larger with larger molecules.
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Noble Gas Molar Mass
(g/mol)
Boiling Point
(K)
He 4.00 4.2
Ne 20.18 27
Ar 39.95 87
Kr 83.80 120
Xe 131.29 165
Dispersion Force and Molar Mass
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-300
-250
-200
-150
-100
-50
0
50
100
150
200
250
1 2 3 4 5 6
Bo
ilin
g P
oin
t, °
C
Period
Relationship Between Dispersion Force and
Molecular Size
BP, Noble Gas
BP, Halogens
BP, XH4
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Practice—The Following Are All Made of Non–Polar Molecules. Pick the Substance in Each
Pair with the Highest Boiling Point.
• CH4 or C3H8
• BF3 or BCl3
• CO2 or CS2
• CH4 or C3H8.
• BF3 or BCl3.
• CO2 or CS2.
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Permanent Dipoles• Because of the kinds of
atoms that are bonded together and their relative positions in the molecule, some molecules have a permanent dipole.Polar molecules.
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Dipole-to-Dipole Attraction
• Polar molecules have a permanent dipole.A + end and a – end.
• The + end of one molecule will be attracted to the – end of another.
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Molar Mass
(g/mol)
Boiling
Point, °C
Dipole
size, D
CH3CH2CH3 44 -42 0
CH3-O-CH3 46 -24 1.3
CH3 - CH=O 44 20.2 2.7
CH3-CN 41 81.6 3.9
Polarity and Dipole-to-Dipole Attraction
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Attractive Forces
+ - + - + - + -
++
++
__
__
Dispersion forces—All molecules.
Dipole-to-dipole forces—Polar molecules.
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Hydrogen Bonding
• Hydrogen atoms bound to a N, O or F atom have strong intermolecular attractions.Unusually high melting and boiling
points.Unusually high solubility in water.
• This kind of attraction is called a hydrogen bond.
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Properties and H-Bonding
Name Formula
Molar
mass
(g/mol)
Structure
Boiling
point,
°C
Melting
point,
°C
Solubility
in water
Ethane C2H6 30.0 -88 -172 Immiscible
Ethanol CH4O 32.0 64.7 -97.8 Miscible
H C
H
H
C H
H
H
H C
H
H
O H
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Intermolecular H-Bonding
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H-Bonds vs. Chemical Bonds
• Hydrogen bonds are not chemical bonds.• Hydrogen bonds are attractive forces
between molecules.• Chemical bonds are attractive forces that
make molecules.
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-200
-150
-100
-50
0
50
100
150
1 2 3 4 5
Bo
ilin
g P
oin
t, °
C
Period
Relationship Between H-Bonding and
Intermolecular Attraction
BP, HX
BP, H2X
BP, H3X
BP, XH4
CH4
NH3
HF
H2O
SiH4
GeH4
SnH4
H2S H2Se
H2Te