States of Matter

States of Matter - Chapter 11

Gases - 11.1

• Kinetic-Molecular Theory (KMT)

  • Explains gas behavior in terms of particles in motion.
  • Assumptions about gas particles:
  • Small size and negligible volume.
  • No significant attractive or repulsive forces between particles.
  • Particles undergo elastic collisions (no kinetic energy loss, energy can be transferred).
  • Temperature correlates with the average kinetic energy (KE) of particles.

• Gas Behaviors

  • Low Density: Gas has fewer particles in a volume compared to solids.
  • Compression: Gases take up large amounts of space and can be squeezed.
  • Expansion: Gases fill their container due to random motion.
  • Diffusion: Movement of one material through another from high to low concentration.
  • Effusion: Gas escapes through small openings.

• Gas Pressure

  • Defined as force per unit area.
  • Units:
  • SI unit: Pascal (Pa)
  • Other units: mm Hg, atm, torr.
  • Conversions: 1 atm = 760 mm Hg = 760 torr = 101.3 kPa.
  • Standard sea level pressure = 101.3 kPa at 0 °C.

• Air Pressure

  • Varies with altitude:
  • Low altitudes: higher pressure, more particles, stronger gravitational pull.
  • Barometers measure atmospheric pressure.

• Atmospheric Pressure Effects

  • At high altitudes, the air is thinner: less pressure and fewer particles available (less O2).
  • Athletes acclimatize to high altitudes before competition.
  • Fewer air resistance effects allow sports equipment (e.g., baseballs, golf balls) to travel farther.

• Dalton’s Law of Partial Pressure

  • Concept is acknowledged but not required for calculations until the gas law chapter.

Forces of Attraction - 11.2

• Intramolecular Forces:

  • Forces holding particles together in bonds. Types include:
  • Ionic: Attraction between cations and anions (Example: NaCl).
  • Covalent: Shared electrons between positive nuclei (Example: H2).
  • Metallic: Mobile electrons among metal cations (Example: Fe).

• Intermolecular Forces:

  • Forces between different types of molecules, weaker than intramolecular forces.
  • Types of intermolecular forces:
  • Dispersion Forces: Result from temporary shifts in electron density; exist between all particles; weakest force.
  • Dipole-Dipole Forces: Attraction between oppositely charged regions of polar molecules; neighboring polar molecules align favorably.
  • Hydrogen Bonds: Special dipole-dipole attraction; occurs with H bonded to N, O, or F.

Liquids and Solids - 11.3

• Liquids

  • Take the shape of their container, fixed volume.
  • Higher density than gases; compressible; can flow (fluid) and diffuse.

• Viscosity:

  • Measures resistance to flow; affected by:
  • Attractive forces: stronger forces = higher viscosity.
  • Particle size and shape: larger/more massive particles = higher viscosity.
  • Temperature: higher temperatures reduce viscosity.

• Surface Tension:

  • Energy needed to increase the surface area of a liquid due to uneven attractive forces.
  • Surfactants (e.g., soaps/detergents) reduce water’s surface tension.

• Cohesion and Adhesion:

  • Cohesion: Attraction between similar molecules.
  • Adhesion: Attraction between different types of molecules.

• Solids:

  • Definite shape and volume; generally denser than liquids and gases.
  • Particles are closely packed; examples include water, where ice floats due to lower density.

• Crystalline Solids:

  • Atoms/ions/molecules arranged in a geometric structure.
  • Unit Cell: Smallest arrangement in a crystal lattice, with seven types differing by angles and face lengths (names not necessary).

• Types of Crystalline Solids:

  • Atomic: Soft, low melting point, poor conductors (e.g., Noble gases).
  • Molecular: Fairly soft, low melting point (e.g., I2, H2O).
  • Covalent Network: Very hard, high melting point (e.g., Diamond, Quartz).
  • Ionic: Hard, brittle, high melting point (e.g., NaCl).
  • Metallic: Soft to hard, malleable, excellent conductors (all metals).

• Amorphous Solids:

  • Irregular particle arrangement; examples: glass, rubber.

Phase Changes - 11.4

• Phase Changes and Energy:

  • Addition/removal of energy causes phase changes.
  • Endothermic Processes: Absorb heat (e.g., melting, sublimation, vaporization).
  • Melting: Solid to liquid; occurs at melting point.
  • Sublimation: Solid to gas (e.g., I2).
  • Vaporization: Liquid to gas, can occur via evaporation.

• Exothermic Processes:

  • Release heat (e.g., freezing, condensation, deposition).
  • Freezing: Liquid to solid at freezing point.
  • Condensation: Gas to liquid.
  • Deposition: Gas to solid.

• Phase Diagrams:

  • Graph of pressure vs. temperature showing phases.
  • Triple Point: Condition where all three phases co-exist.
  • Curving Lines: Indicate conditions where two phases are in equilibrium.