Ch. 10 – Part II Ideal Gas – is an imaginary gas that conforms perfectly to all the assumptions of the kinetic theory. A gas has 5 assumptions 1. Gases consist of large numbers of tiny particles. 2. The particles of a gas are in constant motion, moving rapidly in all directions.

3. The average kinetic energy of the particles of a gas is directly proportional to the temp. of a gas. KE= ½ MV2 4. There are no forces of attraction or repulsion between the particles of a gas. 5. The collision between particles of a gas and between particles of the container’s walls are elastic collisions.

Fluid – a gas or liquid A gas is about 1/1000 the density of the same substance as a solid or liquid. Why? Molecules are farther apart. Compression – gas under pressure. By compressing a gas you can have as much as 100 times more molecules in a cylinder than uncompressed. Effusion – is a process by which gas particles under pressure pass through a very small opening from one container to another. What is diffusion?

Real gas – is a gas that does not behave completely according to the assumptions of the kinetic energy. Johannes van der Waals proposed this. Real gases are explained by the following: 1. Particles of real gases occupy space 2. Particles of real gases exert attractive forces on each other. Gases behave different when heated, cooled, or under pressure. Under “normal conditions” a gas is considered to be an ideal gas.

Gases have 4 measurable quantities 1. Volume 2. Pressure 3. Temperature 4. Quantity of molecules (number) If 3 of these quantities are known then you can figure the fourth one.

If air is heated it can expand its volume many times If air is heated it can expand its volume many times. If it’s cooled it compresses. Pressure is measured by how fast the gas molecules are moving. Determined by how many times the molecules hit the container it is in. Ex. Small vs. large container with 10 molecules. Temp. increases = pressure increase Temp. decreases = pressure decrease Volume decreases = pressure increase Volume increase = pressure decrease

The more molecules of gas in a container, the more pressure it has. Why? In the winter time the pressure in your car tire is less. Why? If you blow up a balloon the volume is constant if the temp. and pressure are constant. IDEAL GAS LAW CONSTANT R = PV/nT

Pressure – is the force per unit area on a surface P = f/a Label  N/cm2 or Pascals 1 N/cm2 = 1 Pascal The SI unit for force is Newtons (N) Barometer – is a device used to measure the atmospheric pressure. Torricelli discovered this. In a vacuum condition and a sea level a colum of mercury or barometer will rise 760 mm.

760 mm of Hg is the atmospheric pressure at sea level and at 0 degrees C. 760 mm of Hg = 760 torr. or 1 atm. of pressure Sample Problem 10.1 Standard conditions or STP – standard temperature and pressure. 1 atm. Of pressure, 760 torr., or 760 mm Hg  Pressure 273 K or 0 degrees Celsius  Temperature 1 Liter or 1000 ml  Volume

Robert Boyle discovered that pressure and volume are inversely proportional to each other. Ex. Double the volume = ½ the pressure Ex. Triple the pressure = 1/3 the volume Ex. Pushing in on the sides of a balloon increases the pressure of the air inside the balloon. P1V1 = P2V2 Sample Problem 10-2

Charles’ Law – states that the volume of a gas varies directly to the temperature of the gas. Ex. Double the temp. = double the volume Ex. Hot air balloon V1/T1 = V2/T2 Sample Problem 10-3

Gay-Lussac’s Law – states that the pressure of a gas is directly proportional to the temp. of the gas. Ex. Double temp. = double the pressure Ex. Car tires P1/T1 = P2/T2 Sample Problem 10-4

Combined gas law – shows the relationship between pressure, volume, and temp. of a gas when the amount of gas is fixed. P1V1/T1 = P2V2/T2 Sample Problems