Chapter 12: States Of Matter

1
Chapter 12 States Of Matter

• Sec. 12.1 Gases

2
Objectives

• Use the kinetic-molecular theory to explain the
  behavior of gases.
• Describe how mass effects rates of diffusion and
  effusion.
• Explain how gas pressure is measured and
  calculate the partial pressure of a gas.

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Properties of Substances

• Chemical physical properties of substances
  depend on composition (the atoms present)
  structure (arrangement of atoms).
• However, substances that are gases display
  similar properties despite different compositions!

4
Kinetic Molecular Theory (1860)

• Gases studied were molecules.
• Objects in motion have energy called kinetic
  energy.
• The kinetic molecular theory describes the
  behavior of gases in terms of particles in motion.

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Kinetic Molecular Theory

• The kinetic molecular theory assumes that gas
  particles have a VERY SMALL volume and that they
  are separated from one another by a LARGE volume
  of space.
• Because they are so far apart, there is no
  attraction or repulsion between gas particles.

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Kinetic Molecular Theory

• Gas particles are in constant, random motion.
• They move in straight lines until collision.
• Collisions between gas particles are elastic.
  (There is no overall loss of kinetic energy.)

7
Kinetic Molecular Theory

• 2 factors determine the kinetic energy of a gas
  particle mass and velocity

Within a gas sample, the mass does not vary but
velocity will. Therefore, when we talk about
KE, we really mean average KE.
8
Kinetic Molecular Theory

• Temperature is a measure of the average kinetic
  energy of the particles in a sample of matter.
• At a given temperature, all gas particles will
  have the SAME average kinetic energy.

9
Behavior of Gases

• The constant motion of gas particles allows a gas
  to expand until it fills its container.

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Behavior of Gases

• Gases have a low density.
• (Remember D mass/volume)
• There are fewer gas particles in a given volume
  than in the same volume of a liquid or solid.
• A great deal of space exists between the gas
  particles.

11
Behavior of Gases

• Gases are compressible (able to have their volume
  reduced) because there is so much empty space
  between gas particles.

12
Behavior of Gases

• Diffusion is the term used to describe the
  movement of one material through another. Gases
  have no forces of attraction for one another so
  diffusion is possible.
• Due to diffusion, gas particles tend to move from
  areas of high concentration to areas of low
  concentration, until they are evenly distributed.

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Behavior of Gases

• Rate of diffusion depends on the mass of the gas
  particles.
• Light particles, at the same temperature as
  heavier particles, will have a greater velocity.
  They will therefore diffuse quicker.
• Effusion is related to diffusion. During
  effusion, a gas escapes through a tiny opening.

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Behavior of Gases

• Grahams law of effusion states that the rate of
  effusion for a gas is inversely proportional to
  the square root of its molar mass. This law can
  also be applied to diffusion rates.

15
Practice Problems

• What is the ratio of the diffusion rate of
  ammonia to hydrogen chloride?
• Calculate the ratio of effusion for neon to
  nitrogen.
• Calculate the ratio of diffusion rates for carbon
  monoxide to carbon dioxide.

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Gas Pressure

• Pressure is defined as force per unit area.
• Gas particles exert pressure when they collide
  with the walls of their container.

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Gas Pressure

• Since pressure is a result of collisions between
  all of the gas particles and the surfaces around
  them, the amount of pressure increases when the
  number of particles in a given volume increases.

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Atmospheric Pressure

• The gas particles in air move in all directions,
  and so, exert air pressure in all directions.
• There is less air pressure at high altitudes
  because there are fewer particles present, since
  the force of gravity is less.
• Torricelli was the first to demonstrate that air
  exerted pressure.
• He invented the barometer, an instrument used to
  measure atmospheric air pressure.

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Atmospheric Pressure

• A barometer has a closed tube that is inverted in
  a pool of Hg. The Hg rises falls in the tube
  in response to the amount of air pressure applied
  to the Hg.

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Atmospheric Pressure

• Torricelli showed that at the Earths surface,
  the height of the Hg in the barometer was always
  about 760 mm Hg. (mm Hg stands for millimeters
  of mercury).
• This is considered standard air pressure.

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Units of Pressure

• The SI unit of pressure is the pascal (Pa).
  Standard air pressure is 101,300 Pa or 101.3 kPa.
• Standard air pressure in more traditional units
  is
• 14.7 psi (pounds per square inch)
• 760 torr (1 torr 1 mm Hg)
• 1 atm (atmosphere)

22
Practice Problems

• Determine the value of each in kPa.
• 3.5 atm
• 930 torr
• 560 mm Hg

23
Measuring Gas Pressure

• A closed or open ended manometer is used to
  measure gas pressure in a closed container.
• In a manometer, the difference in the levels
  levels of Hg in the U-tube is used to calculate
  the gas pressure in mm Hg.

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Daltons Law of Partial Pressure

• This law states that the total pressure of a
  mixture of gases is equal to the sum of the
  pressures of all the gases in the mixture.
• The portion of the total pressure contributed by
  a single gas is called the partial pressure.

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Daltons Law of Partial Pressure

• Mathematically
• P1 P2 P3 . PT
• We add the pressure of each gas in a mixture.
  Their sum is equal to the total pressure of gas
  in the container.

26
Practice Problems

• A mixture of oxygen, carbon dioxide, and nitrogen
  has a total pressure of 0.97 atm. What is the
  partial pressure of oxygen if the partial
  pressure of carbon dioxide is 0.70 atm and that
  of nitrogen is 0.12 atm?
• Find the total pressure for a mixture that
  contains 4 gases with partial pressures of 5.00
  kPa, 4.56 kPa, 3.02 kPa, and 1.20 kPa.