Gases differ from solids and liquids. What aspect of gases, associated with density, do we already know differs from liquids and solids?
I. Properties of Gases A. Gases may be compressed 1. A standard volume of gas may be compressed by adding pressure.
B. Gases expand to fill their container uniformly.
C. All gases have a low density.
D. Gases May be mixed.
II. The Kinetic Theory of Gases and the Ideal Gas Model. A. Gases consist of molecular particles moving at any given instant in straight lines. 1. When an individual particle, moving in a straight line, strikes the wall of its container, it exerts a force.
2. This force reflects a uniform force on all sides of the container and can be deemed “Pressure”.
3. The force of these collisions is directly proportional to the kinetic energy of the particles or the Kelvin temperature of the gas.
B. Molecules collide with each other and with the container without any loss in their kinetic energy.
C. Gas particles are very widely spaced. 1. The ability to compress a gas is directly associated with this large amount of space between the atoms or the molecules. 2. The ability to mix gases is associated with the space between atoms or molecules of a gas.
D. The actual volume of the atoms or the molecules is negligible compared to the space that they occupy. 1. Liquid water at a volume of 1.04 ml., when converted to a gas, occupies a space of 1,670 ml. The original volume of the liquid is a mere 0.6% of the volume occupied by the gas. 2. The volume of the actual molecules in a gas phase is therefore negligible.
E. Gas molecules behave as independent particles, attractive forces between them are negligible.
III. Gas Measurements: Measurements of gases, when the mass of the gas remains constant, deal primarily with the effects of pressure, volume, and temperature.
A. Pressure 1. History: Toricelli: a. The height of a tube of mercury, when placed in a plate of mercury, reflects the atmospheric pressure on the plate of mercury. b. Confirmed by Pascal when he compared heights at different altitudes.
2. Concept of Pressure a. Pressure = force per unit area.
b. English Pounds per square inch.
c. System International (SI) Pascal = one Newton per square meter. More commonly used KPa.
d. Most commonly used: Milliliter of Mercury or Torr and an Atmosphere. When weather stations report barometric pressure they are reporting mm of Hg.
3. Standards of Pressure: One standard atmosphere of pressure is defined as 760 mm of Hg at sea level.
4. Units of Pressure – Relationships
a. 1.0 Atm = 760 mm Hg = 760 Torr b. 1.0 Atm = 1.013 X 105 Pa = 101.3 KPa c. 1.0 Atm = 29.92 in. Hg = 14.69 psi.
5. Boyles Law: Pressure vs. Volume a. Analytical experimentation shows that, with constant gas mass, as pressure increases, volume decreases.
b. Pressure is inversely proportional to volume.
c. Formula: P1V1 = P2V2
d. When three of the four variables are known, the fourth can be calculated.
V2 = P1V1 P2
B. Gas Temperature: Gas Temperatures are usually measured in degree Celsius (deg C). When doing gas problems, however, you must always change to the Kelvin scale.
1. Charles Law: Volume vs.