Chap 10_Liquids_Solids

States of Matter

• States: Three primary states - Gas, Liquid, Solid.

  • Gas:

    • Volume/Shape: Fills the container, no fixed shape.

    • Density: Very low.

    • Compressibility: Very high.

    • Molecular Motion: Vibration & random motion.

  • Liquid:

    • Volume/Shape: Fixed volume, takes shape of container.

    • Density: High.

    • Compressibility: Slight.

    • Molecular Motion: Vibration & translational motion.

  • Solid:

    • Volume/Shape: Fixed volume and shape.

    • Density: High.

    • Compressibility: None.

    • Molecular Motion: Vibration (rigid).

Kinetic Molecular Theory

• Core Concepts: This theory explains behavior of gases.

  • Gases consist of rapidly moving particles in random directions.

  • Gas particles have negligible volume and no mutual attraction.

  • Average kinetic energy (KE) is directly proportional to temperature (in Kelvins).

  • Collisions between gas particles are elastic, keeping total KE constant.

  • Pressure is caused by molecules colliding with container walls.

Intermolecular Forces

• Importance: Determines the state of matter (gas, liquid, solid).

  • At STP, gas particle interactions are minimal; with decreasing temperature/increasing pressure, interactions become significant leading to condensation and solidification.

• Types of Intermolecular Forces:

  • Dipole-Dipole Interactions: Occur between polar molecules, including hydrogen bonds.

  • Ion-Dipole Interactions: Between ions and polar molecules.

  • Van der Waals Forces: Weak forces including London Dispersion Forces (induced dipoles).

  • All molecules exhibit Van der Waals forces, though they are weak compared to intramolecular bonds.

London Dispersion Forces

• Definition: Attraction between temporary induced dipoles.

• Exist between all atoms and nonpolar molecules.

• Strength varies (0.01 to 2 kcal/mol) based on mass, size, and shape of molecules. More contributing factors include:

  • Increased mass and electron count elevate the strength of these forces, despite being weak overall.

Dipole-Dipole Interactions

• Definition: Electrostatic attractions between positive and negative dipoles.

  • Example: Butane (nonpolar) relies solely on London forces; Acetone (polar) involves dipole-dipole interactions.

Hydrogen Bonds

• Definition: Attraction between the partial positive charge of hydrogen (linked to electronegative atoms like O or N) and the partial negative charge on nearby electronegative atoms.

  • Strength varies from 2 to 10 kcal/mol; in water, it is approximately 5.0 kcal/mol.

  • Contributes to high boiling point of water due to extensive intermolecular bonding, requiring additional energy to separate molecules.

Phase Changes

• Phase: Uniform parts of a system (solid, liquid, gas).

• Phase Changes: Transition from one state to another due to thermal energy changes.

• Energy Calculations:

  • Heating 1 gram of solid water from -20°C to 120°C involves calculations using specific heat (SH) for temperature change and latent heat for phase change.

• Phase Diagrams: Graphs showing the state of a substance at varying temperatures and pressures.

Properties of Liquids

• Surface Tension: Result of uneven intermolecular attractions at the surface, creating a thin elastic layer allowing certain objects to float.

• Capillary Action: Ability of a liquid to rise in a narrow tube due to adhesive and cohesive forces.

• Viscosity: Resistance of a fluid to flow, influenced by intermolecular forces and molecular shape.

Properties of Water

• Related to Hydrogen Bonding:

  • High specific heat capacity.

  • High heat of vaporization.

  • Density of ice (solid water) is less than that of liquid water (ice floats).

Solids and Their Structures

• Formation of Solids: Occurs when attractive forces stabilize molecules in a fixed structure through crystallization.

• Types of Solids:

  • Crystalline Solids: Ordered structure, composed of one or more crystals.

  • Amorphous Solids: Disordered structure lacking a well-defined arrangement.

  • Molecular Solids: Held together by intermolecular forces.

  • Metallic Solids: Composed of metal atoms, held by delocalized electrons.

  • Ionic Solids: Formed by cations and anions through ionic bonds.

  • Covalent Network Solids: Atoms connected in large networks by covalent bonds.

Crystal Structures

• Unit Cell: Basic repeating unit in a crystalline structure.

• Coordination Number: Number of neighboring atoms surrounding an atom in a crystal lattice.

• Packing: Different arrangements demonstrate varying strengths and properties of solids.

Conclusion

• Understanding intermolecular forces and phase behaviors is crucial for explaining the physical properties of substances.