1. Gases

2. States of Matter
A. Matter exists in four states: solid, liquid, gas, and plasma.
B. According to the kinetic theory of matter, all matter is made up of tiny particles (atoms, ions, and molecules) in constant motion.
C. It is the motion of the particles, as well as the strength of intermolecular forces in a substance, that determines if it is a solid, liquid, or gas.

3. Characteristics of Gases
Gases have no definite shape nor definite volume.
1. Gas particles of a substance are not held in a fixed position as they are in the solid form of the substance.
2. Gas particles are not held close together by van der Waals forces as they are in the liquid form of the substance.
3. Therefore, gas particles are free to spread far apart from each other, moving randomly in straight lines until they collide with another particle or the wall of its container.

4. Characteristics of Gases
Gas particles have frequent collisions due to very high speeds.
For example, at room temperature (25o C), the average speed of an Oxygen molecule is 433 meters per second. This is equivalent to just over 1700 kilometers per hour.
1. Scientists can determine the average number of collisions a molecule undergoes in a unit of time by finding the average distance a molecule travels before colliding with another molecule.
a. This distance is called the mean free path of the molecule.
2. It is important to note that these factors (the speed, the distance of travel, and the number of molecular collisions) vary with the temperature, the number of particles in a given volume, and the mass of the particles composing the gas.
a. Warmer molecules tend to move faster. Colder molecules move slower.
b. The more in number the more crowded it gets.
c. The smaller the volume, the less room to move around in.
d. Heavier molecules move slower than lighter molecules.

5. Gas Pressure
A. Besides colliding with each other, gas molecules collide with the walls of the container in which the gas is confined.
1. When a gas molecule collides with the wall of a container, it exerts a impact force on the container.
2. It is the force of collision and the number of collisions with the walls of the container that causes gas pressure.
a. The pressure exerted by a gas on its container is the same in every direction.
3. Pressure is measured in terms of the force per unit area (f/area) or pascals.

6. Standard Atmospheric Pressure
A. The molecules and atoms of the gases present in the air (Atmosphere) are constantly hitting the surface of the Earth and everything on it. As a result, everything on Earth’s surface is subject to a certain pressure from the air molecules.
B. Air pressure varies from place to place and from time to time in a particular place.
1. As altitude (rise in height) increases, atmospheric pressure decreases because fewer air particles are found in a given volume.
2. Scientists have agreed on a standard of pressure as representing an average air pressure at sea level.
3. The standard has been set as kilopascals (kPa) and is known as standard atmospheric pressure
4. One pascal is a pressure of one newton per square meter (N/m2).

7. Measuring Pressure
A. In measuring gas pressure, an instrument called a manometer is used.
There are two types of manometers:
1. In an open-arm manometer, the atmosphere exerts pressure on the column of liquid in the open arm of the U-tube as the gas being studied exerts pressure on the other arm.
a. The difference in the liquid level between the two arms is a measure of the difference in pressure between the atmosphere and the contained gas.
b. If the density of the liquid in the manometer is known, then the pressure difference between the gas and the atmosphere can be calculated.

8. Measuring Pressure
2. In a closed-arm manometer, there is a vacuum above the liquid in one arm. The operation of a closed-arm manometer is independent of atmospheric pressure.
a. The difference in the liquid level between the two arms is the pressure of the contained gas.
3. A closed-arm manometer used to measure atmospheric pressure is called a barometer.
a. Most barometers are manufactured with a scale calibrated to read the height of a column of mercury in millimeters.
b. By definition, standard atmospheric pressure will support a column of mercury (Hg) 760mm high.
c. Because standard atmospheric pressure is defined as kPa, it can be stated that kPa = 760 mm Hg.
i. By dividing both sides of the equation by , it is determined that 1 kPa = mm Hg.

9. Gas Pressure

10. Giant Can Crusher

11. Balloon in a Bottle

12. Gas Laws

13. Boyle’s Law
A. The pressure of a gas depends on how often its particles collide with, or strike, the walls of its container.
1. If the gas is squeezed into a smaller space, then its particles will strike the walls more often and gas pressure will increase.
2. Conversely, if the gas particles are given more space, then they will strike the walls less often and the gas pressure will be reduced.
B. Robert Boyle, a British chemist, described this property of gases over 300 years ago. This relationship between pressure and volume of a gas is called Boyle’s Law.

14. Boyle’s Law
If the amount and the temperature of a gas remain constant, the pressure exerted by a gas varies inversely (oppositely) as the volume.
Formula: (pressure1)(volume1) = (pressure2)(volume2)
*or* P1V1 = P2V2

15. Boyle’s Law

16. Gay-Lussac’s Law
A. There is a relationship between the pressure and temperature of a gas at a constant volume.
1. As temperature increases, the kinetic energy of the gas particles increases.
a. The energy and frequency of the collision of gas particles against their containers increases,resulting in an increase in pressure.
B. In the early 1800s, Joseph Louis Gay-Lussac, a French physicist and chemist, announced this relationship between pressure and temperature.

17. Gay-Lussac’s Law
The pressure of a gas increases as the temperature increases, if the volume of the gas is held constant. The pressure decreases as the temperature decreases.
Formula: Pressure of Gas 1 divided by the Temperature of Gas 1 is equal to the pressure of the second gas divided by the temperature of the second gas
*or* P1➗T1=P2➗T2

18. Gay-Lussac’s Law

19. Charles’s Law
A. As a gas is heated, its particles move faster and faster, and its temperature increases.
1. Because the gas particles move faster, they strike the walls of their container more often and with more force.
2. If the walls of the container are free to move, these faster moving gas particles will push the walls out and expand.
B. Jacques Charles, a French physicist, observed the simple relationship between the volume of a gas and temperature over 200 years ago.
1. Charles found that the volume of any gas doubled when the temperature increased from 0oC to 273oC.
2. Charles’s experimental information led to the formation of the absolute, or Kelvin scale.

20. Charles’s Law
The volume of a quantity of gas, held at a constant pressure, varies directly with the temperature
Formula: V1 ➗T1 = V2 ➗T2

21. Charles’s Law

22. Dalton’s Law of Partial Pressure
A. In a mixture of gases, each gas in the mixture contributes to the total pressure. Stated another way, each gas in the mixture exerts its partial pressure.
B. John Dalton, a British chemist, was the first to form a hypothesis about partial pressure.
1. After experimenting with gases, he concluded that each gas exerts the same pressure it would if it were present alone at the same temperature.
2. Dalton’s Law of partial pressure states: The total pressure in a container is the sum of the partial pressures of all the gases in the container.

23. Dalton’s Law of Partial Pressure
Gases in a single container are all at the same temperature and have the same volume. Therefore, the only difference in their partial pressures is due only to the difference in the numbers of molecules present.
Formula: Ptotal= P1 + P2 + P3… Pn

24. Dalton’s Law of Partial Pressure