STATES OF MATTER

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STATES OF MATTER
GASES
LIQUIDS
SOLIDS
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THE NATURE OF GASES
Kinetic theory matter consists of particles
which are in constant motion
Basic assumptions of kinetic theory as applied
to gases

• Gases are composed of particles that have no
  volume
• and are relatively far apart. There are no
  attractive or repulsive
• forces between the particles.

• Gas particles are in constant random motion and
  fill their
• containers regardless of size or shape.

• All collisions between particles are perfectly
  elastic, with
• energy being transferred from one particle to
  another without
• loss. Total kinetic energy remains constant.

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GAS PRESSURE
Gas pressure is defined as the force exerted by a
gas per unit surface area of an object. It is
caused by the gas particles colliding with the
surface of the object.
Vacuum is a space containing no particles and no
pressure.
Atmospheric pressure is the pressure exerted by
the air.
Barometers are devices used to measure
atmospheric pressure.
The SI unit of pressure is the pascal (Pa).
Atmospheric pressure at sea level is about 101.3
kilopascals (kPa). Pressure has also
been measured in mm of mercury and, in the U.S.,
inches of mercury.
One standard atmosphere (atm) is the pressure
required to support 760 mm of mercury at 25oC
1 atm 760 mm Hg 101.3 kPa
Standard temperature and pressure, STP, is 0oC
and 101.3 kPa.
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Kinetic Energy and Kelvin Temperature
Temperature is the measure of the average
kinetic energy of particles in matter.
As a substance is heated, it absorbs energy and
its particles speed up.
This results in an increase in the particles
kinetic energy and an increase in temperature.
The opposite is also true. As a substance cools,
its particles slow and the temperature decreases.
Theoretically, if the temperature is low enough,
all particle motion ceases. This temperature,
absolute zero, is 0 K or 273oC.
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THE NATURE OF LIQUIDS
A MODEL FOR LIQUIDS
1. Liquid particles are in motion.

• The particles can slide past one another (flow)
  and therefore
• will assume the shape of the container in which
  they are held.

• The particles in liquids are attracted to one
  another. The attractive
• force between particles is called intermolecular
  forces.

• In order to change into a gas, the particles in a
  liquid must have
• sufficient kinetic energy to overcome the
  intermolecular forces which
• hold them together.

• The intermolecular forces cause the liquid
  particles to be much
• closer together, more dense. For this reason,
  increasing the pressure
• on a liquid has virtually no effect on its volume.

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EVAPORATION
The conversion of a liquid to a gas or vapor is
called vaporization.
When such a conversion occurs on the surface of a
liquid which is not boiling, the process is
called evaporation.
In evaporation, some molecules in the liquid
break away and enter the vapor state. Only
those molecules of the liquid with a certain
minimum kinetic energy can break away from the
surface. A heated liquid evaporates faster
because the addition of energy increases the
kinetic energy of the particles and allows more
of them to overcome the attractive forces.
Evaporation is a cooling process. The cooling
occurs because the particles with the highest
kinetic energy tend to escape first. The
particles left in the liquid have a lower
average kinetic energy thus the temperature
decreases.
When you perspire, water molecules in your
perspiration absorb heat from your body and
evaporate from your skins surface. The
remaining perspiration is cooler and absorbs more
body heat cooling your body further.
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EVAPORATION IN A CLOSED CONTAINER
When a partially filled container of liquid is
sealed, some of the particles in the liquid
vaporize. These particles collide with the walls
of the container and produce vapor pressure, or
a force due to the gas above the liquid.
Initially, the number of particles entering the
vapor increases but eventually some of the
particles will return to the liquid, or condense.
After a time the number of particles vaporizing
equals the number of particles condensing. At
this point equilibrium occurs, that is, the
rate of vaporization equals the rate of
condensation.
An increase in temperature of a contained liquid
increases the vapor pressure. This happens
because more of the particles have
sufficient kinetic energy to escape as vapor.
The increasing number of particles colliding with
the walls of the vessel results in an increase in
pressure.
Vapor pressure is measured by a device called a
manometer. In a simple manometer, one end of a
U-shaped glass tube containing mercury is
attached to a container. When the container is
empty, the mercury levels on both sides of the
U-tube are equal. When a liquid is added to the
container, the resulting vapor pressure pushes
the mercury on the container side the levels of
mercury are no longer the same
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BOILING POINT
Boiling point is the temperature at which the
vapor pressure of the liquid is just equal to the
external pressure.
If the external pressure changes then the boiling
point of the liquid will change
At sea level (101.3 kPa), the boiling point of
water is 100oC. In Denver, however, water boils
at about 95oC because Denver is 1600 meters
above sea level and has an atmospheric pressure
of only 85.3 kPa.
Boiling is a cooling process similar to
evaporation. The particles with the greatest
energy escape first.
Increasing the heat of a boiling liquid does not
raise its temperature but simply increases the
rate at which the liquid boils.
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THE NATURE OF SOLIDS
A MODEL FOR SOLIDS

• Particles in solids are not free to move and tend
  to vibrate about
• fixed points.

• In most solids the particles are packed against
  one another in
• a highly organized pattern.

3. Solids do not flow to take the shape of their
container.
4. Solids are dense and incompressible

• When you heat a solid the particles vibrate more
  rapidly as their
• kinetic energy increases. The organization of
  particles within the
• solid breaks down and eventually the solid melts.

• The melting point is the temperature at which a
  solid turns into
• a liquid. At this point the disruptive
  vibrations of the particles are
• strong enough to overcome the interactions that
  hold them together.

7. The melting and freezing points of a
substance are the same.
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CRYSTAL STRUCTURE AND UNIT CELLS
Most solid substances are crystalline. A crystal
is made up of atoms, ions, or molecules arranged
in an orderly, repeating, three-dimensional patter
n called the crystal lattice.
The shape of a crystal depends on the arrangement
of the particles within it. The smallest group of
particles within a crystal that retains the
geometric shape of the crystal is known as a
unit cell.
Some solid substances can exist in more than one
form. Allotropes are two or more different
molecular forms of the same element in the same
physical state.
Amorphous solids are solids which lack a
crystalline structure
Glasses are examples of amorphous solids. They
are transparent fusion products of inorganic
substances that have cooled to a rigid state
without crystallizing. Glasses do not melt at a
specific temperature but gradually soften when
heated.
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CHANGES OF STATE
PHASE DIAGRAMS
A phase diagram gives the conditions of
temperature and pressure at which a substance
exists as solid, liquid, and gas (vapor).
The point in a phase diagram at which all three
curves meet is called the triple point. The
triple point describes the only set of
conditions at which all three phases can exist in
equilibrium with one another.
For water, the triple point is a temperature of
0.016oC and a pressure of 0.61 kPa.
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SUBLIMATION
The change of a substance from a solid to a gas
without going through the liquid phase is called
sublimation.
Sublimation occurs because solids, like liquids,
have vapor pressure.
If the vapor pressure is high enough, the solid
will sublimate.
Solid carbon dioxide, dry ice, is an example of a
solid that sublimates. Ice cubes left in the
freezer for a long time become smaller
even though the temperature has never exceeded
the melting point of water.