The document discusses the gas laws and properties of gases. It begins by describing the composition of Earth's atmosphere, which is primarily nitrogen and oxygen. It then discusses that gases have mass and low densities compared to liquids and solids. The document outlines four variables that describe gases - pressure, volume, temperature, and amount. It explains concepts such as gas compressibility, units of measurement for gases, and the kinetic molecular theory which describes gas particles as being in constant random motion.

1. Chemistry
Unit 10:
The Gas Laws

2. The Atmosphere is all around us
 an “ocean” of gases
mixed together
Composition
nitrogen (N2)………….. ~78%
oxygen (O2)…………… ~21%
argon (Ar)……………... ~1%
carbon dioxide (CO2)… ~0.04% Trace amounts of:
He, Ne, Rn, SO2,
water vapor (H2O)……. ~0.1% CH4, NxOx, etc.

3. Elements that exist as gases at 250C and 1 atmosphere

5. Gases have mass.
• Gases seem to be weightless, but
they are classified as matter,
which means they have mass.
• The density of a gas – the mass
per unit of volume – is much
less than the density of a
liquid or solid, however.

6. Gases have mass.
It’s this very low density that allows
us to be able to walk through the
room without concerning ourselves
with air resistance.
Since it is so easy to “swim” across
the room we don’t put much thought
into the mass of a gas.
Really it is only noticeable if we
have a large collection of gas in a
container.

7. Physical Characteristics of Gases
• Gases assume the volume and shape of their containers.
• Gases are the most compressible state of matter.
• Gases will mix evenly and completely when confined to the
same container.
• Gases have much lower densities than liquids and solids.

8. Compressibility
Gases can expand to fill its
container, unlike solids or liquids
The reverse is also true:
They are easily compressed, or
squeezed into a smaller volume
Compressibility is a measure of
how much the volume of matter
decreases under pressure

9. Variables that describe a Gas
There are FOUR variables used to
describe gases:
1. pressure (P)
2. volume (V)
3. temperature (T)
4. amount (n)

10. 1. Pressure of Gas
Pressure results from the collisions
between gas molecules and the walls
of the container they are in. More
molecules means more collisions
which means more pressure.
Gases naturally move from areas of
high pressure to low pressure,
because there is empty space to
move into – a spray can is example.

11. Units of Pressure
The standard unit for pressure is
ATMOSPHERE or atm
 Other units include:
 kPa = kilopascal
 mmHg = millimeter of Mercury
 torr = Torrricelli
 psi = pounds per square inch

12. Convert between Pressure units
1 atm equals...
101.325 kPa
760 mm Hg
760 torr
14.69 psi

13. 2. Volume of Gas
In a smaller container, the
molecules have less room to move.
The particles hit the sides of the
container more often.
As volume decreases, pressure
increases. (think of a syringe)
Thus, volume and pressure are
inversely related to each other

14. Units of Volume
The standard unit for
volume is the...
LITER or L
*There are 1,000 mL in 1 L

15. 3. Temperature of Gas
 Temperature is a measurement of the
amount of Kinetic Energy the gas molecules
contain
 Raising the temperature of a gas increases
the pressure, if the volume is held constant.
(Temp. and Pres. are directly related)
The molecules hit the walls harder, and
more frequently!
 Should you throw an aerosol can into a fire?
What could happen?

16. Units of Temperature
The standard unit for temperature
is...
Kelvin or K
Other units include:
Degrees Celsius = oC
Degrees Fahrenheit = oF

17. Absolute Zero
The theoretical temperature at
which all kinetic motion
completely stops. Equal to 0 K
or -273 oC
Conversions
K = °C + 273
°C = K – 273

18. 4. Amount of Gas
When we inflate a balloon, we are
adding gas molecules.
Increasing the number of gas
particles increases the number of
collisions
thus, the pressure increases
 The standard unit for the amount
of gas molecules is the: MOLE (mol)

19. And now, we pause for this
commercial message from STP
OK, so it’s really not THIS kind
of STP…
STP in chemistry stands for
Standard Temperature and
Pressure
Standard Pressure = STP allows us to compare
1 atm amounts of gases between
different pressures and
Standard Temperature temperatures
= 0 oC or 273 K

20. Kinetic Molecular Theory
The theory states that the tiny
particles in all forms of matter in
all forms of matter are in
constant motion.
This theory is used to explain
the behaviors common among
gases
There are 3 basic assumptions
of the KMT as it applies to gases.

21. Kinetic Molecular Theory of
Gases
Three basic assumptions of the kinetic
theory as it applies to gases:
#1. A gas is composed of small,
particles that have mass- usually
molecules or atoms. They have...
Insignificant volume; relatively far apart
from each other
No attraction or repulsion between
particles

22. Kinetic Molecular Theory of
Gases
#2. Particles in a gas move rapidly
in constant random motion
Move in straight paths, changing
direction only when colliding with one
another or other objects
Average speed of O2 in air at 20 oC is
an amazing 1700 km/h!

23. Kinetic Molecular Theory of
Gases
#3. Collisions are perfectly
elastic- meaning kinetic energy is
transferred without loss from one
particle to another- the total
kinetic energy remains constant

24. Summary of The Kinetic Molecular Theory
-- explains why gases behave as they do
1. …are so small that they are
assumed to have zero volume
2. …are in constant, straight-line motion
3. …experience elastic collisions in which
no energy is lost
4. …have no attractive or repulsive forces toward
each other
5.…have an average kinetic energy (KE) that is
proportional to the absolute temp. of gas
(i.e., Kelvin temp.) as Temp. , KE

25. Kinetic Molecular Theory of
Gases
Gas Pressure – defined as the force
exerted by a gas per unit surface
area of an object
Due to: a) force of collisions, and b)
number of collisions
No particles present? Then there cannot
be any collisions, and thus no pressure
– called a vacuum

26. Kinetic Molecular Theory of
Gases
Atmospheric pressure results from
the collisions of air molecules with
objects
Decreases as you climb a mountain
because the air layer thins out, meaning
less particles, as elevation increases
Barometer is the measuring device
for atmospheric pressure, which is
dependent upon weather & altitude

27. Atmospheric pressure changes with altitude:
As altitude , pressure .
barometer: device to measure
air pressure
vacuum
mercury air
(Hg) pressure

28. Measuring Pressure
The first device for
measuring atmospheric
pressure was developed
by Evangelista Torricelli
during the 17th century.
The device was called a
“barometer”
Baro = weight
Meter = measure Torricelli

29. An Early
Barometer
The normal pressure
due to the atmosphere
at sea level can
support a column of
mercury that is 760 mm
high.

30. Kinetic Molecular Theory of
Gases
Whathappens when a substance is
heated? Particles absorb energy!
Some of the energy is stored within the
particles- this is potential energy, and
does not raise the temperature
Remaining energy speeds up the
particles (increases average kinetic
energy)- thus increases temperature

31. Kinetic Molecular Theory of
Gases
Anincrease in the average kinetic
energy of particles causes the
temperature to rise.
As it cools, the particles tend to move
more slowly, and the average K.E.
declines.
Is there a point where they slow down
enough to stop moving?

32. Kinetic Molecular Theory of
Gases
Theparticles would have no kinetic
energy at that point, because they
would have no motion
Absolute zero (0 K, or –273
oC) is the
temperature at which the motion of
particles theoretically ceases
This has never been reached, but about
0.5 x 10 -9 K has been achieved

33. •Diffusion:
describes the mixing
of gases. The rate of
diffusion is the rate of
gas mixing.
•Molecules move
from areas of high
concentration to low
concentration.

34. Effusion: a gas escapes through a tiny
hole in its container
-Think of a nail in your car tire…