General Properties of Liquids
and Types of Solids
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How Important Liquids Is?
• Liquids are vital to our lives.
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Water is a means of food preparation
Cooling machines n industrial processes
Recreation
Cleaning
Transportation
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What makes this possible?
• Floating needle/paper clip
• Water strider floating
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CHARACTERISTICS OF LIQUIDS
• Surface tension
• Capillary action/Capillarity
• Viscosity
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SURFACE TENSION
• A measure of the
inward forces that
must be overcome in
order to expand the
surface area of a
liquid.
• The greater the forces
of attraction between
molecules of the
liquid, the greater the
surface tension.
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Surface Tension Cont.
• Surface tension of a
liquid decreases with
increasing
temperature.
• The stronger the
intermolecular forces
the stronger the
surface tension.
Water has a high surface
tension do to hydrogen
bonding. 6
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What makes this possible?
• Why do liquids rise through a narrow tube
such as the capillary tube?
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CAPILLARY ACTION
• Another way surface
tension manifests.
• The rise of liquids up
very narrow tubes.
This is limited by
adhesive and
cohesive forces.
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Capillarity
• The cohesive forces between the like
molecules, that s, the IMF within the liquid,
compete with adhesive forces of unlike
molecules or the forces between the liquid
and the walls of the capillary tube.
• Capillarity is the ability of liquids to rise in a
narrow tube because the adhesive forces are
greater than the cohesive forces.
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COHESIVE FORCES
• Intermolecular forces that bind like
molecules to one another (e.g. hydrogen
bonding).
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ADHESIVE FORCES
• Intermolecular forces that bind a
substance to a surface.
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Formation of meniscus
• Water : adhesive
forces are greater
than cohesive forces
• Mercury: Cohesive
are greater than
adhesive forces.
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VISCOSITY
• The resistance of a liquid to flow.
• The less “tangled” a molecule is expected
to be, the less viscous it is.
Water = less Viscosity
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syrup = high Viscosity
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Viscosity Cont.
• Viscosity decreases with increasing
temperature (molecules gain kinetic
energy and can more easily overcome
forces of attraction).
• Viscosity Increases as pressure increases.
• Liquids with strong IMF have a higher
viscosity.
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Structure of Solids
• Two ways to categorize solids
– Crystalline
– Amorphous
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Properties of Solids
• Solids may be classified according to their
structure. Those that have a well-defined
shape due to the orderly arrangement of
their atoms, molecules or ions are called
CRYSTALLINE SOLIDS.
• Those which are disorganized are called
AMORPHOUS SOLIDS.
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Crystalline Solid
• Ridged and long range order of its atoms.
• Solids have flat surfaces
• Sharp/High melting points
• EX: Quartz, diamond, sodium Chloride.
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Crystalline Solids
• A crystalline solid has a well-defined
crystal lattice.
• Lattice- is a 3dimensional system of points
designating the positions of the components
(atoms, ions, molecule) that make up a crystal.
• Unit cell- smallest repeating unit of a lattice.
There are seven common unit cells and lattices.
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Common Unit Cells and Lattices
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Amorphous Solid
•
•
•
•
Lack a well defined arrangement
No long range order
IMF vary in strength
DO NOT have sharp melting points.
EX: rubber, glass
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FORCE/S
HOLDING THE
UNITS
TOGETHER
GENERAL
PROPERTIES
EXAMPLES
Electrostatic
attraction
Hard, brittle, high
melting point,
poor conductor
of heat/electricity
NaCl, LiF,
MgO,CaCO3
Covalent
Covalent
Hard, high
melting point,
poor conductor f
heat/electricity
C (Diamond),
SiO2 (quartz)
Molecular
Dispersion
forces, dipoledipole forces,
hydrogen bonds
TYPES OF
CRYSTAL
Ionic
Metallic
Metallic bond
Soft, low melting
Ar, CO2, I2,
point, poor
H20,Cl2H22O11,
conductor of
(sucrose)
heat/electrcity
Soft to hard, low
to melting point,
good conductor
All metallic
elements
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Changes of state
• Transformation from one state to another
Condensation
Vaporization
AKA: steam
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Changes in state
• Liquid  Gas Vaporization Endothermic
• Gas  Liquid Condensation Exothermic
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• Solid  Gas
Sublimation
Endothermic
• Gas  Solid
Deposition
Exothermic
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• Solid  Liquid Melting
Endothermic
• Liquid  Solid Freezing
Exothermic
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Changes of state
• The energy involved it phase changes is
calculated using
– Heat of fusion (solid  liquid or liquid solid)
– Heat of vaporization (liquid gas or gas liquid)
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Vapor Pressure (vp)
Vapor Pressure: Pressure
exerted by molecules that have
enough energy to escape the
surface.
As T ↑ VP ↑evaporation ↑
Liquids with high VP are volatile
(alcohol evaporates easily)
Liquids that have strong IMF have
low vapor pressures.
(take a lot of energy to overcome
IMF so it can evaporate)
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How to make something boil
1. Increase the VP of the liquid (heat it) so
that the VP of the liquid is > that of the
atmosphere.
2. Lower the atmospheric pressure
(pressure above the liquid)
3. At high altitudes (low air pressure) water
boils at a lower temperature
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Boiling Point
The vapor pressure of the liquid = external atmospheric pressure
Note: The normal boiling point of water is 100oC. The term
normal refers to standard pressure or 1 atm, or also 30
101.3
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kPa.
Boiling Point
Increasing the temperature of a liquid increases the
kinetic energy of its molecules until such point where the
energy of the particle movement exceeds the IMF that
hold them together.
NOTE: The boiling point of a liquid is influenced by the
strength of its IMF. The greater the attractive forces, the higher
the energy needed to increase the KE of the molecules to
break these forces.
At higher altitudes, the atmospheric pressure is lower; hence,
the boiling point will subsequently decrease.
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Boiling Pts. of H2O at Various
Elevations
Altitude compared
to Sea Level
(m)
Boiling
Point
(°C)
1609
98.3
177
100.3
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Heat of Vaporization
• Molar Heat of vaporization- is the
amount of heat required to vaporize one
mole of a substance at its boiling point.
The application of heat disrupts the IMF of
liquid molecules and allows them to
vaporize.
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