Chemistry gases

• States of Matter
• Gases


  • Shape: Not fixed (fills a container)
  • Volume: Not fixed (fills space)
  • Particle motion: Constant random motion
  • Arrangement: Particles are very far apart

  
• Gases do not all behave the same way
• Kinetic Molecular Theory of Gases


  A theory that explains the behavior of gases
• Assumptions of the Kinetic Molecular Theory

  • Gases consist of small particles separated from one another by empty space
  • Gas particles are in constant, random motion
  • Collisions between gas particles are elastic
  • There are no significant attractive or repulsive forces between gas particles
  • Particle size is negligible compared to the distances between particles

  
• Kinetic Energy (KE)


  The energy of motion, depends on mass and velocity

  
• Temperature


  A measure of the average kinetic energy of the particles in a sample of matter

  
• Gases have low density compared to solids and liquids
• Compression and Expansion of Gases

  Changing the volume occupied by a constant mass of gas particles
• Diffusion


  The movement of one material through another
• Effusion


  A gas escaping through a tiny opening
• At the same temperature

  Gases have the same average kinetic energy
• Graham's Law


  The rate of effusion or diffusion of a gas is inversely proportional to the square root of its molar mass

  
• Molar masses

  • Ammonia (NH3): 17.0 g/mol
  • Hydrogen chloride (HCl): 36.5 g/mol
• The ratio of diffusion rates of ammonia and hydrogen chloride is 1.47
• Barometer

  An instrument that measures air pressure
• Manometer


  An instrument that measures gas pressure in general
• Dalton's Law of Partial Pressures

  The total pressure of a mixture of gases is equal to the sum of the partial pressures of all the gases in the mixture
• Partial pressure can be used to calculate the amount of gas produced in a chemical reaction