Ch. 13 States of Matter

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Ch. 13 States of Matter

• Ch. 13.1 The Nature of Gases

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Kinetic Theory and a Model for Gases

• Kinetic energythe energy an object has because
  of its motion.
• Assumptions about gases
• The particles in a gas are considered to be small
  hard spheres with an insignificant volume.
• The motion of the particles in a gas is rapid,
  constant, and random.

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• All collisions between particles in a gas are
  perfectly elastic (no energy is lost).

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3 states of matter solid, liquid, gas
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Motion in gases vs solids

• gases solids

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Gas Pressure

• Gas pressure results from the force exerted by a
  gas per unit surface area of an object.
• Gas pressure is the result of simultaneous
  collisions of billions of rapidly moving
  particles in a gas with an object.
• Vacuuman empty space with no particles and no
  pressure.

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• Atmospheric Pressureresults from collisions of
  atoms and molecules in air with objects.
  Decreases with increasing elevation.
• Barometers measure atmospheric pressure. Normal
  atmospheric pressure at sea level can push
  mercury up a vacuum tube to a height of 760 mm.

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• Pascal (Pa)the SI unit of pressure. Very small
  unit. Normal atm. pressure is about 100,000 Pa.
• One standard atmosphere (1
  atm) 760 mm Hg 101.3 kPa at 25C.
• Pressure units conversion problems

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Kinetic Energy and Temperature

• Higher temperature of a substance means there is
  a higher average speed, or kinetic energy, of its
  atoms or molecules.
• At any given temperature, all substances,
  regardless of state, have the same average
  kinetic energy.
• Temperature is directly related to the average
  kinetic energy of a substances particles.

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• Absolute zerono particle motion no kinetic
  energy. 0 K. -273.15C. Never produced in the
  laboratory.
• The Kelvin temperature of a substance is directly
  proportional to the average kinetic energy of the
  substance.