Gases and the Kinetic-Molecular Theory. Chapter 12. Common Properties of Gases. Gases can be compressed into smaller volumes by applying increased pressure Gases exert pressure on their surroundings Gases expand without limits Occupy the volume of any container
Note: Pressure in Rexburg is less. Why?
Pgas(torr) = Patm(torr) + h torr
Pgas(torr)=Patm(torr) - h torr
Problems: On chalkboard
Illustration: Pressure apparatus
Lord Kelvin noticed that the extension of the volume to 0, produced a value of –273.15C. This was defined as absolute zero on the Kelvin scale. Using this scale, a quantitative relationship was established.
Demo: Balloon with liquid nitrogen
This serves as a reference point for discussing gases.
For a given sample of gas . From
this relationship, . This is the combined gas law.
Demo: Expanding a balloon (assume P and T constant)
at constant T and P
What would be the volume of 0.25 moles of O2 gas be at STP?
Remember, the ideal gas law is used to describe one set of conditions.
A 1.502-gram sample of a pure gaseous compound occupies 852 mL at 34.2C and 845 Torr. What is the molecular weight?
Look at page 456 for the derivation of Dalton’s Law
The total pressure in a sealed container containing water is affected by the pressence of water. Gaseous molecules leaving the liquid phase produces a contribution to the total pressure. This contribution is called the vapor pressure of water. The vapor pressure of water increases with increasing temperature (Table 12-4). Actually, every liquid has a unique vapor pressure.
Any gas in contact with water soon becomes saturated with water vapor.
Therefore, each gas exerts a partial pressure, and the total pressure is a sum of all the molecule-wall collisions.
It should be noted, however, that the urms is not equal to the diffusion rate of a gas through air. Why?
How will Videallyavailable vary with He and C3H8?
The Van der Waals equation reduces to the ideal gas law at high temperatures and low pressures. Why?