1. Chemistry
Properties of Gases

2. Properties of Gases
Gases have unique properties because the distance between the particles of a gas is much greater than the distance between particles in liquids and solids.
Gases are considered fluids because they are able to flow. (Don’t confuse this with liquids, which are also able to flow!)

3. Kinetic Molecular Theory explains the properties of gases
Gas particles are in constant rapid, random motion.
The particles of a gas are very far apart relative to their size.
Gas particles collide with each other and with the walls of their container.

4. Kinetic-Molecular Theory—one of science’s most successful theories!
4. The pressure exerted by a gas is a result of collisions of the molecules against the sides of the container.
5. The kinetic-molecular theory considers collisions of gas particles to be perfectly elastic—this is energy is transferred completely during collisions.

5. Gases Have Low Density
Gases have low density—much less than liquids or solids.
Because of the large distances between gas particles, most of the volume of a gas is empty space.
Gas particles travel long distances before colliding with each other.

6. Gases are Highly Compressible
The space occupied by the gas particles themselves is very small compared with the total volume of gas.
Applying a small amount of pressure will move the gas particles closer together and will decrease the volume.
A bicycle pump works by compressing air into the tire using the action of a piston inside the pump

7. Gases Completely Fill a Container
Gas particles are far apart and moving independently of each other at high speed.
They will quickly fill any container they are put into because they are in constant, random, straight-line motion.

8. Gas Pressure
Earth’s atmosphere is a mixture of gases: mainly nitrogen and oxygen.
Because you cannot always feel air, you may have thought of gases as weightless.
But, all gases have mass; therefore they have weight in a gravitational field.

9. Air Pressure
As gas molecules in the atmosphere are pulled toward the surface of Earth due to gravity, they collide with each other and with the surface of the Earth more often. These collisions are what causes air pressure.

10. Atmosphere (Unit of Pressure)
Atmospheric air pressure can be measured by a barometer.
The atmosphere exerts pressure on the surface of mercury in the dish.
This pressure goes through the fluid and up the column of mercury.
The mercury settles at a point where the pressure exerted downward by its weight equals the pressure exerted by the atmosphere

11. Evangelista Torricelli (1608-1647)
The barometer was invented by the Italian physicist and mathematician Evangelista Torricelli in 1643.
Torricelli was a student of Galileo.
The pressure unit torr was named in his honor.

12. may be subject to copyright.
Atmospheres and Torrs
At sea level, the atmosphere keeps the mercury in a barometer at an average height of 760 mm, which is 1 atmosphere.
One millimeter of mercury is also called a torr.
As we gain altitude by going up a mountain, the air pressure decreases. Why?

13. Atmospheric Air Pressure and Storms
Changes in atmospheric air pressure often signal changes in weather patterns.
Low pressure air systems signal storms and winds.
Where is a storm
system located on
this map? →

14. Measuring Pressure Pressure is: force/area
To find pressure, you need to know the force and the area over which that force is exerted.
The unit of force in SI terms is called the newton (N)
The SI unit of pressure is called the pascal (Pa)
The pascal is named for Blaise Pascal, one of the founders of modern physics.

15. Units of Pressure 1 Pa = 1 N/1 m2
One Pascal is a small unit of pressure. It is the pressure exerted by a layer of water that is only mm deep over an area of 1 square meter.
A pressure gauge showing psi (pounds/square inch) in red and kPa (kiloPascals) in black

16. Atmospheric pressure can be expressed a number of ways:
1 kiloPascal = 1,000 Pa
1 atmosphere (atm) = 101,325 Pa
1 bar (bar)= 100,025 Pa
1 Millimeter of Mercury (mm Hg) = Pa
Pounds per square inch (psi) 1 psi= X 103 Pa
1 Torr (torr) = Pa

17. The Gas Laws
In this section, you will study the relationship between the measurable properties of a gas, represented as follows:
P=pressure exerted by the gas
T=temperature in Kelvins of the gas
V=total volume of the gas
n=number of moles of the gas

18. Pressure-Volume Relationship—Boyle’s Law
Robert Boyle, an English scientist, studied the relationship between volume and pressure in 1662.
If the temperature and the amount of gas are kept constant, when the volume of a gas is decreased, the pressure increases.
The product of the
pressure and volume
(PV) remains almost
constant if the
temperature stays the
same.

19. Boyle’s Law: If volume ↓, pressure ↑
There is an inverse relationship between pressure and volume.

20. Boyle’s Law: P1V1=P2V2
If the temperature and the number of particles of the gas are not changed, then the PV product remains the same. Therefore, P1V1=P2V2

21. Temperature-Volume Relationships: Charles’s Law
In 1787, the French physicist Jacques Charles discovered that a gas’s volume is directly proportional to the temperature in Kelvins.
Kelvin= degrees Celsius + 273

22. Charles’s Law: If temp ↑, volume ↑
Heating a gas makes it expand. Cooling a gas makes it contract.
The direct relationship between temperature and volume is known as Charles’s Law.

23. Charles’s Law: If Temp ↓,Volume ↓
If all other conditions are kept constant, then Charles’s Law can be expressed as follows:

24. Temperature—Pressure Relationships: Gay-Lussac’s Law
Pressure is the result of collisions of particles with the walls of the container.
If the temperature of the gas particles is doubled, their average kinetic energy (energy of motion) is also doubled.
If there is a fixed amount of gas in a container of fixed volume, the collisions will have twice the energy so the pressure will double.

25. Gay-Lussac’s Law: Temp ↑ Pressure ↑
The French chemist Joseph Gay-Lussac is given credit for discovering this law in 1802.
There is a direct positive relationship between temperature and pressure.

26. Gay-Lussac’s Law: Temp ↓ Pressure ↓
If the volume of the container is constant and there is an equal amount of gas particles, when temperature ↑, then pressure ↑. Gay-Lussac’s Law can also be stated as:

27. Avogadro’s Law
Amedeo Avogadro determined that equal volumes of gases (V) at the same temperature (T) and pressure (P) will have the same number of gas particles (n).
Avogadro’s Law states that
there is a direct relationship
between the number of gas
particles (n) and the volume.
If (n) increases, then V increases.

28. Avogadro’s Law
If the gas particles change from one type of gas to another but the number of gas particles is the
same, they will
have the same
volume.
Avogadro’s
Formula:

29. An Ideal Gas
You have examined several laws related to gases. No gas perfectly obeys all of these laws under all conditions. One way to model a gas’s behavior is to assume that the gas is an ideal gas that perfectly follows these rules.
An ideal gas:
1. does not condense to a liquid at low temperatures
2. does not have forces of attraction or repulsion between the particles
3. is composed of particles that have no volume.

30. The Ideal Gas Law (also called the Combined Gas Law)
All of the previous examples assume that two properties are held constant and one property is changed to see what happens.
However, life is rarely that simple. We can combine these laws into one equation!
R = joules per kelvin per mole (J · K-1 · mol-1).

31. Ideal Gas Law The Ideal Gas Law means that:
For a constant P, T increases, V increases.
For a constant V, T increases, P increases.
For a constant T, P increases, V decreases.