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BACK                                                               Gases

I  A Model to Explain Gas Behavior
    A.  the nature of gases
        1.  all gases exhibit similar physical behavior
              a.  gases have mass
              b.  easy to compress
              c.  fill their containers completely
              d.  different gases can move through each other easily (diffusion)
              e.  exert pressure
               f.  gas pressure depends on temperature
        2.  properties explained by the Kinetic-molecular model
    B.  the kinetic-molecular model
        1.  developed by Rudolf Clausius, James Maxwell,and Ludwig Boltzman
        2.  postulates
              a.  gas consist a very small particles that have mass
              b.  distance separating gas particles are large
              c.  in constant rapid, random motion
              d.  collisions are perfectly elastic
              e.  average kinetic energy of gases depend on their temperature
              f.  gas particles exert no force on one another

II  Measuring Gases
    A.  gas behavior dependent on four variables:
        1.  amount of gas (n)
              a.   measured in moles
              b.  n   =
    mass     
                         molar mass
        2.  volume (V)
              a.  gas volume = container volume
              b.  measured in liters(dm3)
        3.  temperature (T)
              a.  measured in degrees Celsius
              b.  calculations must use the Kelvin scale
              c.  T(K) = T(
0C) + 273
        4.  pressure (P)
              a.  particles colliding with container walls exerting a force
              b.  force spread over the area is called pressure
    B.  atmospheric pressure and the barometer
        1.  pressure measured with a barometer
        2.  atmospheric pressure - mass of air attracted by gravity
        3.  units:
              a.  1 atmosphere (atm) = 101,325 Pa
              b.  1 atm = 760 mm Hg = 760 torr
              c.  1 atm = 14.70 lb/in.
2
              d.  1 bar = 100,000 Pa = .9869 atm
    C.  enclosed gases
        1.  pressure measured using a manometer
        2.  U-shaped tube filled with mercury - connected to a container of gas the other endopen to the atmosphere
        3.  if the two ends of mercury are equal heights - atm pressure = gas  pressure
        4.  if mercury level lower on the gas side of tube - gas pressure greater than the atm
        5.  if level higher on the gas side of tube - gas pressure less than atm
        6.  to find gas pressure
              a.  if gas side lower - add difference in level heights to the atm
              b.  if level higher - subtract difference in level heights from the atm pressure
    D.  STP
        1.  comparison of gases are done at standard temperature and pressure
        2.  standard temperature = 0oC
        3.  standard pressure = 1 atm


III  The Gas Laws
    A.  Boyle's law: the pressure-volume relationship
        1.  performed a series of experiments
        2.  used a J-shaped tube closed at one end
        3.  air was trapped by pouring mercury in the tube
        4.  measured the volume of air when different amounts of mercury were added
        5.  multiplied the volume times the pressure = same product  PV= k1
        6.  indicated a relationship between pressure and volume
        7.  if  the pressure increases the volume must decrease in order for the product to be constant, k1
        8.  states the pressure and volume of a sample of gas at constant temperature are inversely proportional to each other
        9.  allows the calculation of changes in the volume of a gas when the pressure changes
                        P1V1 = P2V2
                            V2 =
P1V1
                                       P2
    B.  Charles' law:  the temperature-volume relationship
        1.  developed the relationship between temperature and volume
        2.  used a cylinder with a movable top and water at various temperatures
        3.  plotted volume against temperature - shows a straight line plot
        4.  indicates volume is directly proportional to its temperature
        5.  lowest possible temperature for gases = -273.15
0C
        6.   -273.15
0C known as absolute zero
        7.  absolute temperature scale or Kelvin scale has its zero as absolute zero
        8.  Charles' law states that at constant pressure, the volume of a fixed amount  of gas is directly proportional to its absolute temperature
        9.  allows the calculation of changes in the volume of a gas when the temperature changes
                     V1T1 = V2T2
                         V2 =
V1T1
                                     T2
    C.  Avogadro's law:  the amount-volume relationship
        1.  states that equal volumes of gases at the same temperature and pressure contain an equal number of particles
        2.  two important points:
               a.  all gases show the same physical properties
               b.  larger volume - greater number of particles
        3.  1mol = 22.4 L at STP called molar volume
    D.  Dalton's law of partial pressure
        1.  different gases in a mixture act independently
        2.  each gas exerts its own pressure as if it was alone
        3.  each pressure called partial pressure of the total pressure
        4.  states that the sum of the partial pressures of all the components in a gas mixture is equal to the total pressure of the gas mixture
        5.  PT = p
a + pb + pc+...
             PT
=  total pressure
                          p
a + pb + pc +...  partial pressures


IV  The Ideal Gas Law
    A.  the ideal gas equation
        1.  PV=nRt
        2.  describes the physical behavior of an ideal gas
        3.  relates pressure (P), temperature (T),volume (V), and moles (n)
        4.  ideal gas - described by the kinetic-molecular theory postulates
        5.  no actual gas exists
        6.  does describe real gases under normal condition
        7.  does not at very low temperatures and high pressures
        8.  R - the ideal gas constant (universal)
              a.  units:
  1.  .0821 atm-L / mol-K
  2.  8.314 Pa-m3 / mol-K
  3.  8.314 J / mol-K
              b.  measured units must be the same as the constant units
    B.  ideal gas law and the kinetic-molecular theory
        1.  equation (V-T constant): P increases with Increase in mol theory:  increase particles increase collisions, therefore pressure
        2.  equation (n-V constant):  increase temperature - increases pressure theory:  increase in kinetic energy increases force of the collisions, therefore pressure
        3.  equation (n-T constant):  decrease volume - pressure increases theory:  particles in smaller volume increases number of collisions, therefore pressure
        4.  experimental data (ideal gas equation) supports kinetic-molecular theory
    C.  deviations from ideal behavior
        1.  high pressure
              a.  decreases volume of the gas
              b.  particle volume begins to be a significant portion of the overall gas volume
              c.  ideal equation begins to fail
              d.  ideal equation assumes particles have no volume
        2.  low temperature
              a.  reduces kinetic energy
              b.  attractive forces between particles become significant
              c.  ideal equation begins to fail
              d.  ideal equation assumes no attractive force between particles
        3.  ideal gas equation still useful provided you understand its limitations


V.  How Gases Work
    A.  lifting power of gases
        1. gases less dense than air will rise
        2.  density can be modified using the principles of the ideal gas equation
    B.  gas effusion
        1.  diffusion gas particles move through another
        2.  effusion
              a.  related to diffusion
              b.  particles pass through a hole one particle at a time
        3.  lighter particles diffuse and effuse faster than heavier gases





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