# Solids and Liquids

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```					     States of Matter
“It is the city of mirrors, the city of
mirages, at once solid and liquid, at
once air and stone.”
Erica Jong
All assumptions about matter come
from Kinetic Molecular Theory
• All molecules of a pure substance are identical.
• Motion of molecules obeys Newton’s Laws, but
– Molecules move randomly
– They move in any direction with equal probability
– Distribution of speeds does not change over time (certain % hi
speed, medium speed, low speed)
• Collisions (particle-particle and particle-wall) are all
elastic, therefore Kinetic Energy is conserved.
•   Negligible forces exist between molecules.
•   # of molecules is large, but the average separation
between molecules is great compared with their
dimension.
•   Volume of molecules themselves is negligible compared
to volume of container.
WHAT DO YOU THINK?

• Are these assumptions reasonable?
• Get together in groups of 3. Examine each
assumption. Tell me if you think that each
of these assumptions are reasonable. Why
or why not? Be specific!
Three common phases of matter
Three Phase of Matter
• Solid –

• Liquid –

• Gas –
Forces and Phases
• Substances with very little intermolecular
attraction exist as gases
•   Substances with strong intermolecular
attraction exist as liquids
•   Substances with very strong
intermolecular (or ionic) attraction exist
as solids
Relative Magnitudes of Forces

• The types of bonding forces vary in
their strength as measured by
average bond energy.
Strongest
• Covalent bonds (400 kcal/mol)
• Hydrogen bonding (12-16 kcal/mol )
• Dipole-dipole interactions (2-0.5
kcal/mol)
•   London forces (less than 1 kcal/mol)

Weakest
Wait, what is pressure?
• Pressure = Force/area
• That is what allows this to happen:
Gas Pressure
• An empty region of
space is called a
_________.
• Gas Pressure inside a
vacuum is ______
•   A __________ is used
to measure gas
pressure
How do we measure pressure?
• Atmospheric pressure
– collisions of air
molecules with
objects.
•   1 atm = 760 mm Hg
= 760 torr =
101,325 Pascal =
101.3 kPa
•   1Pascal = 1N/m2
Effect of Altitude on Pressure
Atmospheric Pressure at Various Locations
Location            Feet above         Patm
sea level        (kPa)
Top of Mt Everest, Tibet     29,028            32
Top of Mt. Denali, Alaska    20,320            45.3
Top of Mt. Whitney,        14,494            57.3
California
Top of Mt. Washington,        6,293            78.6
NH
Boulder, Colorado          5,430            81.3
Madison, Wisconsin           900             97.3
New York City, NY           10             101.3
Death Valley, California      -282            102.3
And what is temperature?

• We know that it is a measure of energy.
• Kinetic Energy = ½ m v2
• Mass is constant for a particle, so
temperature must be a measure of _____
of a particle.
So how do we measure it?
• The scale we use is the Kelvin scale. It is a
direct measurement of the speed of molecules.
•   When particle motion stops, temperature is zero
Kelvin
•   This is called ______________
•   K = °C + 273.15
•   Melting point of H2O = 273.15 K
•   Boiling point of H2O = 373.15 K
Intermolecular Forces in Liquids
• Surface Tension
– A force that tends to pull adjacent of a liquid’s surface
together, thereby decreasing surface area to the
smallest possible size.
Intermolecular Forces in Liquids
• Causes particles to attract and bond while still
moving
•   Empty space between particles is reduced,
therefore liquids have a greater density
– Exception: Water is one of only a few substances that
is less dense as a solid that as a liquid.
• Increasing pressure has little effect on liquids,
since there is not much space to compress
Intermolecular forces in liquids

• Viscosity
– The resistance to ______ that is exhibited by
all liquids and gases.
– The greater the viscosity the ______ the
substance flows.
– Oil is ______ viscous than water (even
though it has lower density!)
– Temperature will effect viscosity. How?
Liquid to gas
• Evaporation/Vaporization

• Molecules at the surface of the liquid break away and go
into gas state.
•   Must have sufficient kinetic energy to overcome surface
tension.
Liquid to Gas

• Evaporation
– Liquids evaporate ________ when heated
because their increased kinetic energy
overcomes the attractive forces at the surface
– Liquids evaporate _______ if they have higher
intermolecular forces
– Volatility: tendency for a liquid to go to a gas
at a given temperature
Evaporation in a closed container
• Particles vaporize, escaping into the
atmosphere above the liquid.
• They collide with the walls of the
container and each other, and condense
back into the liquid.
Evaporation in a close container
• Dynamic equilibrium - the rate of
condensation of gas _________ the rate
of evaporation of liquid.
• In general, the higher the temperature,
the _________ the vapor pressure.
Evaporation
Which liquid has the highest intermolecular forces? The
lowest?
Evaporation is a cooling process
• Particles with the ________ kinetic energy
(temperature) escape first.
• Particles with the ________ kinetic energy
(temperature) remain behind.
• When the “hottest” particles are removed,
the average temperature of the remaining
molecules is lower.
• Examples: Sweating, canteens.
Boiling point and vapor pressure

• High boiling point is a sign of ______
intermolecular forces (and low vapor
pressure).
• _______ temperatures are needed to
break IM forces, which keep particles from
vaporizing.
Vaporization

• Boiling Point
– The temperature at which the vapor pressure
of a liquid equals the pressure exerted on the
liquid.
– Heating allows particles at the liquid’s surface
to break the attractive forces keeping them in
the liquid state.
– Remaining particles gain energy, until all
particles have enough energy to vaporize.
Boiling Point
• At higher elevations, boiling point _________
because particles leaving the liquid are less likely
to collide with external air molecules or other
vapor molecules, therefore they need _____
energy to escape the liquid.
•   At higher elevations: boiling occurs at lower
temperature (energy) so food needs more time
to obtain necessary energy to cook.
•   Pressure cookers reduce cooking time by
_______ pressure, and therefore cooking
temperature.
Effect of Pressure on Boiling Point
Boiling Point of Water at Various Locations
Location            Feet above       Patm      Boiling
sea level      (kPa)      Point oC
Top of Mt Everest, Tibet     29,028          32          70
Top of Mt. Denali, Alaska    20,320         45.3         79
Top of Mt. Whitney,        14,494         57.3         85
California
Top of Mt. Washington,        6,293         78.6         93
NH
Boulder, Colorado          5,430         81.3         94
Madison, Wisconsin           900          97.3         99
New York City, NY            10         101.3         100
Death Valley, California       -282        102.3        100.3
Evaporation vs Boiling
• Evaporation is a surface phenomenon - since the vapor
pressure is low, and the pressure inside the liquid is
equal to atmospheric pressure plus the liquid pressure,
bubbles of water vapor cannot form. But at the boiling
point, the saturated vapor pressure is equal to
atmospheric pressure, bubbles form, and the
vaporization becomes a volume phenomena.
•
The Nature of Solids

• Particles are tightly packed, in an
organized pattern.
• Dense, incompressible, don’t flow.
• When solids are heated, the particles
_______________ as energy increases.
• The structure of particles crumbles and
the solid melts.
Melting Point

• The temperature when a solid turns to a
liquid.
• Particles have enough energy to overcome
some interactions.
• Ordered structure is destroyed.
• Same as freezing point
Melting Point

• Ionic solids have ______ melting points
because they are held together by strong
IM forces.
• Covalent substances have ______ melting
points.
• Not all solids melt. Can you name one?
What is this process called?
Sublimation

• When a solid goes from a solid to a gas
(skips the liquid phase)
• CO2 (dry ice) sublimes at -78˚C
• Solids have a vapor pressure, but usually
it is very low.
• Examples: Dry ice, iodine, snow, freeze-
dried coffee
Types of Solids

• Crystalline Solids: highly regular
arrangement of their components.
• NaCl
Unit Cell
The smallest portion of a crystal lattice that
shows the three-dimensional pattern of
the entire lattice
Glass structure
• Glasses have amorphous structure (no
long range pattern)
• Glass is technically a liquid that has
infinitely low flow.
Phase Diagram
• Represents phases as a function of
temperature and pressure.
•   Critical temperature: temperature above
which the vapor can not be liquefied.
•   Critical pressure: pressure required to
liquefy AT the critical temperature.
•   Critical point: critical temperature and
pressure (for water, Tc = 374°C and 218
atm).
•   Triple point: The only conditions at which
all three states exist in equilibrium.
Phase diagram

Pressure
atm                   Liquid
Solid
Critical Point

Triple Point

Gas

Temperature     oC
Phase diagram of water
Phase transitions
Solid  Liquid
Solid  Gas
Liquid  Solid
Liquid  Gas
Gas  Liquid

Gas  Solid
Phase diagram of water
Phase diagram of water

• Water’s solid liquid line slants backwards,
telling us that when you increase the
pressure, you ________ the melting point
– Solid ice is less dense than liquid ice, which is
why ice floats
– Ice skaters depend on this for creating a film
of water for skating on.
Phase diagram for CO2
Phase diagram for Carbon

Graphite and diamond are Allotropes, 2 or more forms of
the same element in the same physical state.
Plasma
• A mix of electrons and positive ions (ionized
gas)
•   Most common phase of matter in the visible
universe
•   Atoms are moving fast
•   Occurs if we heat a gas to a very high
temperature
•   Partial plasmas – neon lights, aurora borealis,
etc.

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 views: 4 posted: 8/3/2011 language: English pages: 43