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.