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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 MV

- 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 and container walls are elastic collisions.

- Fluid – a gas or liquid directly proportional to the temp. of a gas.
- 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 according to the assumptions of the kinetic energy.

- 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. according to the assumptions of the kinetic energy.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 it has. Why?
- 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

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