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KINETIC – MOLECULAR THEORY OF GASES

KINETIC – MOLECULAR THEORY OF GASES. The kinetic molecular theory is based on the idea that particles of matter are in constant motion. Gases consist of molecules whose separation is much larger than the size of the molecules themselves. Most of the volume of a gas is empty space .

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KINETIC – MOLECULAR THEORY OF GASES

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  1. KINETIC – MOLECULAR THEORY OF GASES The kinetic molecular theory is based on the idea that particles of matter are in constant motion. Gases consist of molecules whose separation is much larger than the size of the molecules themselves. Most of the volume of a gas is empty space. Particles in a gas constantly move in straight line paths and random directions. Particles in a gas collide frequently with the sides of the container and less frequently with each other. All collisions are elastic (energy may be transferred between particles, but no energy is gained or lost as a result of the collisions). Particles in a gas do not attract or repel one another. The amount of kinetic energy in the gas is directly related to the temperature of the gas. (Meaning that as kinetic energy increases, temperature increases.) All gases at the same temperature have the same amount of kinetic energy.

  2. You need to remember…… These postulates describe ideal gases. Real gases act like ideal gases at normal temperatures and pressures. However at low temperatures or high pressures, the behavior of real gases is significantly different than ideal gases. Low Temperatures High Pressures

  3. DEFINITIONS,VARIABLES, AND THEIR UNITS Gas - phase of matter with no definite shape or volume Kinetic Energy-energy of motion Pressure -the force exerted by a gas per unit area on a surface Partial Pressure- pressure of one gas in a mixture of gases Temperature- (Kelvin) directly related to the amount of kinetic energy of a substance STP- Standard Temperature and Pressure = 1.0 atm and 0.0˚C.

  4. DEFINITIONS,VARIABLES, AND THEIR UNITS phase of matter with no definite shape or volume GAS – KineticEnergy – Pressure – PartialPressure– Temperature – STP – Energy of motion • the force exerted by a gas per unit area on a surface pressure of one gas in a mixture of gases (Kelvin) directly related to the amount of kinetic energy of a substance Standard Temperature and Pressure = 1.0 atm and 0.0˚C.

  5. Conversions: *The temperature MUST be in KELVIN UNITS for any calculation in this unit! Remember:   Example: *It doesn’t matter what unit Pressure and Volume are in, as long as you use the SAME UNIT within a single calculation. Remember: Examples: K = oC + 273 1 atm = 760 mm Hg = 101.3 kPa

  6. Combined Gas Law: Example Calculation: What would be the volume in liters of an 8.90 liter sample of gas at 100.oC and 113 kPa if conditions were changed to STP? DON’T FORGET: The term STP stands for Standard Temperature and Pressure. Mathematically, this means 1.0 atmosphere pressure (or 760 mm Hg or 101.3 kPa) and 0.0oC (or 273 K).

  7. Boyle’s Law Increase Pressure = Decrease Volume Decrease Pressure = Increase Volume Real life application:When you breathe, your diaphragm moves downward, increasing the volume of the lungs. This causes the pressure inside the lungs to be less than the outside pressure so air rushes in. Example Calculation:1.00 L of a gas at standard temperature and pressure is compressed to 473 mL. What is the new pressure of the gas?

  8. Charles’s Law Increase Temp = Increase Volume Decrease Temp = Decrease Volume Real life application:Bread dough rises because yeast produces carbon dioxide. When placed in the oven, the heat causes the gas to expand, and the bread rises even further. Example Calculation:A man heats a balloon in the oven. If the balloon initially has a volume of 0.4 liters and a temperature of 20 0C, what will the volume of the balloon be after he heats it to a temperature of 250 0C?

  9. Gay-Lussac’s Law Increase Temperature = Increase Pressure Decrease Temperature = Decrease Pressure Real life application:The air pressure inside a tire increases on a hot summer day. Example Calculation: If the pressure on a gas in a rigid container changes from 800. mmHg to 2.15 atm, and the temperature is found to be 40.0°C, what was the original temperature in Kelvin?

  10. Avogadro’s Law Increase number moles = Increase Volume Decrease number moles = Decrease Volume Real Life Application: Adding air to a balloon (adding more moles) will increase the size (volume) of the balloon. Example Calculation: A 36.2 mole sample of carbon dioxide occupies a volume of 5.20 liters. If 15.4 more liters are added, how many moles of gas would you have if the pressure and temperature remain constant?

  11. Dalton’s Law of Partial Pressure: Ptotal = PA + PB + PC + …for however many gases you have! What if the units are not the same??? Example Calculation: There is a container which has oxygen, xenon and helium in it. Its total pressure is known to be 972 mmHg. If the pressure of the helium is 0.458 atm and the pressure of the oxygen is 74.1 kPa, what is the pressure of the Xenon in mmHg?

  12. Dalton’s Law Over Water: Example Calculation: 50.0 ml of hydrogen collected over water and has a pressure of 850. mmHg at 27.0oC. What is the pressure of the dry gas at STP? The gas collected includes the desired gas AND water vapor… so… Ptotal = Pwater + Pgas (torr = mmHg)

  13. Ideal Gas Law: P= V= n= R= T= PV = nRT Example Calculation: How many moles of oxygen must be put in a 0.500 L flask at 20.00C to have a pressure in the flask of 1.77 atm?

  14. Example Calculation: If 1.17 grams of helium are put into a 500.0 ml container at 100C, what would be the pressure of the container (in mmHg)?

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