Cram AP Chem Unit 3: Intermolecular Forces and Properties

Introduction to Unit 3: Intermolecular Forces and Gas Laws

• Overview of topics covered:

  • Intermolecular forces (hydrogen bonding, dipole-dipole, London dispersion)

  • Ideal gas law and deviations

  • Spectroscopy and Beer's Law

Intermolecular Forces

• Definition:

  • Forces of attraction between molecules, crucial for understanding physical properties.

• Types of intermolecular forces:

  1. London Dispersion Forces:

     • Occurs between non-polar molecules and noble gases.

     • Weak attractive forces due to temporary dipoles.

  2. Dipole-Dipole Interactions:

     • Occurs between polar molecules.

     • Stronger than London dispersion; involve permanent dipoles.

  3. Hydrogen Bonding:

     • Special type of dipole-dipole interaction with hydrogen bonded to N, O, or F.

     • Strongest intermolecular force.

  4. Ion-Dipole Forces:

     • Occur between ions and polar molecules (strongest overall).

London Dispersion Forces

• Characteristics:

  • Formed due to temporary dipoles in non-polar molecules.

  • Dependent on polarizability (size and electron count) and contact area.

• Example comparisons:

  • CCl4 vs. CH4: CCl4 is larger and has stronger London dispersion forces.

  • C2H6 vs. C4H10: Butane has more electrons and a larger contact area, resulting in stronger London dispersion forces.

Dipole-Dipole Interactions

• Characteristics:

  • Occur between molecules that are polar.

  • Example: HCl aligns with oppositely charged regions attracting one another.

• Important note: All polar molecules exhibit both dipole-dipole and London dispersion forces.

Hydrogen Bonding

• Description:

  • Strongest intermolecular force; occurs in highly polar molecules with H-N, H-O, or H-F bonds.

• Example: Water's H bonds create unique properties like high boiling point and surface tension.

Ion-Dipole Forces

• Definition:

  • Attractive forces between ions and polar molecules, crucial for solubility.

• Example: NaCl in water forms ion-dipole interactions with cations and anions surrounded by water molecules.

Ideal Gas Law

• Formula:

  • PV = nRT,

    • P = pressure, V = volume, n = moles, R = gas constant, T = temperature.

• Characteristics of gases:

  • Gas pressure is related to collision frequency; more particles = higher pressure.

• Example application: Reducing volume in a sealed container raises pressure.

Spectroscopy and Beer's Law

• Definition:

  • Study of how light interacts with matter, used to analyze concentrations of solutions.

• Spectrophotometry process:

  • Light source -> monochromator -> sample solution -> detector.

• Beer's Law:

  • A = abc,

    • A = absorbance, a = absorptivity, b = path length, c = concentration.

• Relationship:

  • Higher concentration = higher absorbance; linear relationship.

Key Examples and Applications

• Distillation: Used for separating liquids with different boiling points (e.g., alcohol and water).

• Chromatography:

  • Paper chromatography separates components based on polarity and size.

  • Column chromatography allows substances to move through a stationary phase and separate.

• Electromagnetic Spectrum:

  • Range includes gamma rays, x-rays, visible light, infrared, microwaves, radio waves.

  • Key relationships: Higher frequency (Nu) = shorter wavelength (Lambda).

• Photoelectric Effect:

  • Phenomenon demonstrating the particle nature of light; photon energy = E = hNu.

  • Illustrates electron ejection based on photon energy (dependent on frequency).

Types of Solids

Ionic Solids

• Composed of ions held together by electrostatic forces (ionic bonds).

• Typically have high melting and boiling points.

• Non-conductive in solid form but conductive when melted or dissolved in water.

• Example: Sodium chloride (NaCl).

Molecular Solids

• Composed of molecules held together by intermolecular forces (such as Van der Waals forces, hydrogen bonds).

• Generally have lower melting and boiling points.

• Poor conductors of electricity and heat.

• Example: Ice (solid water).

Metallic Solids

• Made up of metal atoms bonded by metallic bonds, characterized by a sea of delocalized electrons.

• Exhibit high electrical and thermal conductivity.

• Malleable and ductile, with varying melting points.

• Example: Copper (Cu).

Covalent Network Solids

• Atoms connected by covalent bonds in a continuous network.

• Extremely hard with very high melting points.

• Poor conductors of electricity; typically insulators.

• Example: Diamond (a form of carbon).