Kinetic Molecular Theory and
the Gas Laws
Phases of Matter, Kinetic Molecular Theory,
and Temperature
Gas Laws
Phase Changes
Phases of Matter
There are four phases of matter:
1. Solid
2. Liquid
3. Gas
4. Plasma
The state of matter depends on the motion of the
molecules that make it up.
Solids
Solids are objects that have definite shapes and
volumes. The atoms or molecules are tightly
packed, so the solid keeps its shape. The
arrangement of particles in a solid are in a regular,
repeating pattern called a crystal.
Microscopic picture of a solid.
Liquids
The particles in a liquid are close together, but are
able to move around more freely than in a solid.
Liquids have no definite shape and take on the
shape of the container that they are in.
Microscopic picture of a
liquid.
Gases
A gas does not have a definite shape or volume. The
particles of a gas have much more energy than
either solids or liquids and can move around
freely.
Microscopic picture of a gas.
Plasma
Plasma is a gas-like mixture of positively and
negatively charged particles. Plasma is the most
commonly found element in the universe, making
up 99% of all matter. It is found in stars, such as
the sun, and in fluorescent lighting. Plasma occurs
when temperatures are high enough to cause
particles to collide violently and be ripped apart
into charged particles.
Kinetic Molecular Theory of Matter
All matter is made of particles that are in
constant motion. The more energy the
particles have, the more freely they move
around. This freedom that molecules have
is the determining factor for their state of
matter. Therefore, solids have the least
amount of energy. Liquids come next
followed by gases. Finally, plasma has the
most energy of any state of matter.
Temperature
Temperature is a measure of the amount of the
average kinetic energy of the particles in matter.
The more kinetic energy the particles have, the
higher the temperature. The temperature of
particles are usually recorded in one of three ways:
1. Fahrenheit (ºF)
2. Celsius (ºC)
3. Kelvin (K)
Do you remember which is the standard unit????
Fahrenheit
Developed by Daniel Gabriel Fahrenheit, who is best
known for inventing the alcohol thermometer and
mercury thermometer in the early 1700’s. It is based
on 32º for the freezing point of water and 212º for
the boiling point of water. The interval between the
freezing and boiling points are divided into 180 parts.
The conversion to Celsius is:
ºF = (9/5 ºC) + 32
Celsius
Scale developed by Anders Celsius in the early to
mid-1700’s, working from the invention of
Fahrenheit's thermometers. The Celsius scale is
based on 0º for the freezing point of water and
100º as the boiling point. The interval between the
freezing and boiling points are divided into 100
parts.
The conversion to Fahrenheit is: ºC= (5/9)(ºF-32)
The conversion to Kelvin is: K=ºC +273
Kelvin
Developed by William Thompson Kelvin in 1848,
Kelvin is a temperature scale having an absolute
zero below which temperatures do not exist. At
0K, all molecules cease any type of motion (as in
the temperature of outer space). It corresponds to a
temperature of -273° on the Celsius temperature
scale. The Kelvin degree is the same size as the
Celsius degree, so the freezing point of water is at
273K and the boiling point is at 373K.
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The Behavior of Gases
The behavior of gases can be explained by the way
their particles interact with each other and the
environment around them.
The particles are constantly colliding with one
another and other objects. Since the molecules
have mass, there is a certain amount of pressure
being applied.
As the volume of the gas and/or the temperature of
the gas change, so does its behavior.
Gas Laws
The result of a force distributed
over an area.
SI unit for pressure = pascal (Pa) =
N/m2
(one kilopascal = kPa= 1000 Pa)
Factors that Affect Pressure of an
Enclosed Gas
Temperature
Volume
Number of Particles
Temperature
Raising the temperature of a gas
will increase its pressure if the
volume of the gas and the number
of particles are constant
Volume
Reducing the volume of a gas
increase its pressure if the
temperature of the gas and the
number of particles are constant.
Number of Particles
Increasing the number of particles
will increase the pressure of a gas
if the temperature and the volume
are constant.
Boyle’s Law
Boyle’s Law shows the relationship between the
volume and pressure of a gas:
The volume of a fixed amount of gas varies
inversely with the pressure on the gas.
If the pressure increases, the volume decreases; if the
pressure decreases, then the volume increases.
P1V1=P2V2
Boyle’s Law
From Physical Science, Merrill, 1993
Boyle’s Law
A Graph of Boyle’s Law
Charles’s Law
Charles’s Law shows the relationship between the
temperature and volume of a gas:
The volume of a fixed amount of gas varies
directly with the temperature of the gas.
If the temperature increases, the volume increases; if the
temperature decreases, then the volume decreases.
V1 = V2
T1 T2
A Graph of Charles’s Law
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The Combined Gas Law
The relationships described by Boyle’s law and Charles’ law can be
described as a single law. The combined gas law describes the
relationship among the temperature, volume, and pressure of a gas when
the number of particles is constant.
P1V1 P2V2

T1
T2
Phase Changes
A reversible physical change that occurs when a
substance changes from one state of matter to
another.
The temperature of a substance doesn’t change
during a phase change.
Energy is either absorbed or released during a
phase change.
Heat of fusion = energy a substance must absorb
in order to change from a solid to a liquid.
Six Common Phase Changes
Melting- temperature at which a substance
changes from solid to liquid.
2. Freezing – temperature at which a
substance changes from a liquid to a solid.
3. Evaporation – substance changes from a
liquid to a gas. (Heat of Vaporizationenergy a substance must absorb in order to
change from a liquid to a gas.)
1.
4. Condensation- substance changes from a
gas or vapor to a liquid.
5. Sublimation – substance changes from a
solid to a gas or vapor without changing to a
liquid first (endothermic)
6. Deposition – substance changes directly
into a solid without first changing to a
liquid (exothermic)
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