Kinetic Molecular Theory and Gas Laws Day 1

Kinetic-Molecular Theory – explains how particles in matter behave 1.All matter is composed of small particles that are far apart. Gas is mostly empty space. 2. Particles are in constant, random motion. 3.Particles collide with each other and walls of their containers: collisions create pressure 4. Collisions are elastic = no KE lost 5. No attractive/repulsive forces between particles. Molecules move in straight lines.

PHASESOLIDLIQUIDGAS Kinetic EnergyLowMediumHigh ShapeRigidFluidFluid – fills container; diffusion ForcesStrongMediumWeak VolumeDefinite Indefinite CompressibilityIncompressible COMPRESSIBLE! “squish!” DensityHigh (particles close together) MediumLow (spread out)

EXIT QUESTIONS: 1) Each of these flasks contains the same number of molecules. Which container has the highest pressure? Explain your answer.

EXIT QUESTIONS: 2)Which of the following changes to a system will NOT result in an increase in pressure? Explain why you chose your answer. a)Increasing the volume of container b)adding more gas molecules c)Decreasing the volume of the container d)Raising the temperature

Kinetic Molecular Theory and Gas Laws Day 2

Factors affecting gases 1.Volume – amount of space an object occupies Measured in milliliters (mL) or Liters (L) 1000 mL = 1 L We already have heard that 1 mol = 22.4 STP The more moles we have the bigger the balloon will need to be !

Example 1 How many moles of nitrogen gas are in 89.6 L at STP? 89.6 L N 2 1 mol N L N 2 = 4.00 mol N 2

Example 2: What volume does 76 grams of fluorine (F 2 ) occupy at STP (normal conditions)? 76 g F g F 2 1 mole F L F 2 = 44.8 L F 2 **45 L

2. Temperature  Average kinetic energy of particles (how fast they go)  Measured in Kelvin  K = o C Ex: Convert 17 o C to Kelvin: 17 o C = 290 K

*C= 5/9 (*F-32) *F= 9/5 (*C) + 32 K= *C *C= K- 273

3. Pressure Force exerted by a gas per unit area on a surface. Example: Pounds/in 2 or psi Results from the simultaneous collisions of billions of gas particles with the walls of the vessel containing the gas. Standard pressure: 760 mm Hg = 1 atmosphere = kPa = in. Hg = 14.7 psi = 760 torr

Measuring Atmospheric Pressure Measured with a barometer. A barometer uses a column of mercury that rises to an average height of 760 mmHg at sea level. 1 atmosphere (1 atm)

Standard Temperature and Pressure (STP) The conditions of standard temperature and pressure are = 1.0 atm pressure and = 273 K (or 0  1 mole of gas = 22.4 L of gas

Example 1 The atmospheric pressure in Denver, CO is atm on average. Express this pressure in mm Hg.

0.83 atm  mm Hg 1 atm = 760 mm Hg 0.83atm 1 atm 760 mm Hg = mm Hg **631 mmHg

Example 2 Convert a pressure of 175 kPa to atmospheres.

175 kPa  atm kPa = 1 atm 175 kPa kPa 1 atm 1.72 atm

Gas Law Foldable Fold the left and right to the middle. Cut along solid lines (but only to the crack!)

Dalton’s Law of Partial Pressure The pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases. P Total = P 1 + P 2 + P 3 ….

Example A balloon is filled with air (O 2, CO 2, & N 2 ) at a pressure of 1.3 atm. If P O 2 = 0.4 atm and P CO 2 = 0.3 atm, what is the partial pressure of the nitrogen gas?

P Total = P 1 + P 2 + P 3 …. P total = P O 2 + P CO 2 + P N atm = 0.4 atm atm + P N 2 P N 2 = 0.6 atm