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Chapter 13 States of Matter. Read pgs. 384 - 409. Kinetic Molecular Theory. The kinetic molecular theory describes the behavior of gases in terms of particles in motion. A. Gases consist of point-masses. No volume B. Gas particles are in constant motion.

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Chapter 13 States of Matter

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chapter 13 states of matter

Chapter 13States of Matter

Read pgs. 384 - 409

kinetic molecular theory
Kinetic Molecular Theory
  • The kinetic molecular theory describes the behavior of gases in terms of particles in motion.
    • A. Gases consist of point-masses. No volume
    • B. Gas particles are in constant motion.
    • C. Collisions between gas particles are elastic.
    • D. There are no repulsive or attractive forces acting between gas particles.
    • E. Temperature is a measure of the average speed of gas particles.
behavior of gases
Behavior of Gases
  • Gases have low densities. Gases have lots of empty space between particles.
  • Gases have no definite volume. They can be compressed or expanded.
  • Gases diffuse evenly through out a space based upon their temperature and molar mass.
  • Kinetic Energy = ½ m v2 where:
  • m = mass and v = velocity, speed
graham s law of effusion
Graham’s Law of Effusion
  • At the same temperature, two gases will effuse at a rate based upon their molar mass.
  • The heavier the gas the slower it moves.
  • The lighter the gas the faster it moves.
  • See board for equation.
gas pressure
Gas Pressure
  • Because gas has mass, it creates a pressure. Pressure = Force / Area
  • The higher up in the atmosphere you go, the less air, the less air pressure.
  • Air pressure is the weight of air over an object.
  • Units used to measure pressure: Atmospheres, (Atm), Kilopascals, (KPa), millimeters of mercury, (mmHg)

Standard Pressure is pressure at sea-level at 0oC.

  • 1.0 atm = 101.3 kPa = 760 mm Hg
  • Dalton’s Law of Partial Pressures

Ptotal = P1 + P2 + P3 + . . . + Pn

The total pressure in a closed container is equal to the sum of all pressures in the container.

intermolecular forces
Intermolecular Forces
  • Intermolecular forces are forces that hold particles together. They are not the same as a chemical bond.
  • Dispersion Forces – weak forces caused by the movement of electrons.
  • Dipole Forces – attractive forces between polar molecules.
  • Hydrogen Bonds – special type of dipole force between hydrogen and an unshared pair of electrons.
properties of liquids
Properties of Liquids
  • The kinetic molecular theory can also be applied to liquids and solids.
  • Liquids are more dense than gases because the particles are closer together.
  • Liquids are only slightly compressible.
  • Liquids have no definite form so they can flow. Liquids are fluid.
  • Liquids have a viscosity. Viscosity is a measurement of the resistance to flow.

Viscosity is temperature dependent.

  • The higher the temperature, the more a liquid flows.
  • Think pancake syrup.
  • Cold syrup – pours slowly. Warm syrup – pours very quickly.
  • Surface Tension – is a measurement of the inward pull by particles on the surface of a liquid.
  • The more surface tension a liquid has, the more it will ball up.
  • Think about water on a freshly waxed car.
  • Raindrops are round because of this.
properties of solids
Properties of Solids
  • Solids vibrate about fixed points.
  • Solids tend to be the densest phase of matter for most substances.
  • Two types of solids: Crystalline and Amorphous
  • Crystalline solids have a definite arrangement of atoms.
    • Crystals come in definite shapes
    • Crystalline solids have definite melting points

Amorphous solids has no regular repeating pattern.

  • Amorphous solids have no definite melting point.
  • You can also classify solids by the type of bonds that they have.
    • Molecular solids – covalent bonds
    • Ionic solids – ionic bonds
    • Metallic solids – metallic bonds
phase changes
Phase Changes
  • All phase changes are reversible.
  • All phase changes comes in pairs.
    • Freezing – Melting solids and liquids
    • Vaporization – Condensation liquids and gases
    • Sublimation – Deposition solids and gases
  • All phase changes involve a change in energy. One is exothermic and one is endothermic.
  • For a phase change to occur, particles must have some minimum kinetic energy.
melting freezing
Melting - Freezing
  • Melting points and freezing points are always the same.
  • When particles have enough energy to overcome Intermolecular forces, phase change happens.
  • Temperature remains constant during phase change.
vaporization evaporation
Vaporization / Evaporation
  • Vaporization is the change of a liquid to a gas.
  • Vaporization or Boiling take place at a specific point. Evaporation can take place at any temperature.
  • Boiling points vary, but take place through out the whole liquid.
  • Evaporation only happens at the surface of a liquid.
  • Evaporation is a slow change from a liquid to a gas.
  • Evaporation can be increased by several ways.
    • Increase the room temperature.
    • Increase air currents over liquid.
    • Increase surface area of liquid.

Evaporation can even take place at temperatures below freezing.

boiling points
Boiling Points
  • For boiling to take place, particles have to have a certain amount of energy.
  • Boiling point is also affected by atmospheric pressure.
  • Normal boiling point is the boiling point at standard pressure.
  • Water can boil at any temperature.
sublimation deposition
Sublimation / Deposition
  • Sublimation is the direct change from the solid phase to the gaseous phase.
  • Deposition is the opposite of sublimation.
  • Some examples of sublimation:
    • Dry ice, frozen CO2
    • Freezer burn
phase diagrams
Phase Diagrams
  • Phase diagrams show the phases of a substance at different temperatures and pressures.
  • Each substances phase diagram is a little different.
The triple point is the exact temperature and pressure where all three phases exist at the same time.
  • The critical point is where you can no longer change a gas back to a liquid.