AP Chemistry Unit 3 Review: Properties of Substances and Mixtures

Intermolecular Forces

• London Dispersion Forces

  • Weakest intermolecular forces, but can be stronger in large molecules.
  • Only force present in nonpolar molecules.
  • More electrons → more polarizable → stronger forces.

• Dipole-Dipole Forces

  • Present in polar molecules; positive pole of one attracts negative pole of neighbor.
  • Generally stronger than dispersion forces.

• Hydrogen Bonding

  • Occurs in molecules with H bonded to O, F, or N.
  • Especially strong intermolecular force.

• Ion-Dipole Forces

  • Between polar molecules and ions in ionic compounds.
  • Important for dissolving ionic compounds in polar solvents like water.

• Force Strength Order: London < Dipole-Dipole < Hydrogen

Types of Solids and Their Properties

• Ionic Solids

  • High melting points, brittleness, electrical conductance when dissolved.

• Covalent Network Solids

  • Examples: Diamond, Silicon Dioxide.
  • Strong covalent bonds, each atom bonded in multiple directions.

• Molecular Solids

  • Individual molecules as units (e.g., sugar), weak forces, low melting points.

• Metallic Solids

  • Pure metals and alloys, malleable, ductile, good conductivity.

• True Solids vs. Amorphous Solids

  • Crystalline structure vs. non-crystalline (e.g., plastics).

States of Matter

• Solid: Particles close, vibrational motion.
• Liquid: Particles farther apart, can flow.
• Gas: Particles independent, far apart, compressible.

Gas Laws and Behavior

• Ideal Gas Law: PV=nRT

  • P = Pressure (atm), V = Volume (L), n = Moles, T = Temperature (K), R = 0.08206 L atm/mol K.

• Partial Pressures

  • Total pressure = sum of partial pressures of gases.
  • Partial pressure = mole fraction x total pressure.

• Graphical Relationships

  • Pressure, Volume, Temperature, Moles correlations.

• Temperature and Kinetic Energy

  • Same temperature = same average kinetic energy.
  • Higher temperature = higher average kinetic energy.

• Boltzmann Distribution

  • Visualizes molecule speed distribution at different temperatures.

• Ideal Gas Conditions

  • High temperature, low pressure for minimal particle interaction.

Mixtures and Solutions

• Heterogeneous vs. Homogeneous Mixtures (Solutions)

• Molarity

  • Moles of solute / Liters of solution.

• Diagrams

  • Mole ratios, states of components (solid, aqueous).

Separation Techniques

• Distillation

  • Different boiling points to separate components.

• Chromatography

  • Separation based on adherence to stationary phase.

Solubility

• "Like Dissolves Like"
  • Polar dissolves polar, nonpolar dissolves nonpolar.

Electromagnetic Spectrum and Molecular Effects

• Ultraviolet/Visible Light: Electron transitions.
• Infrared Radiation: Molecular vibrations.
• Microwave Radiation: Molecular rotations.

Light and Photons

• Equations:
  • c = λν (speed of light, wavelength, frequency).
  • E = hν (energy of photon, Planck's constant).

Spectrophotometry and Beer-Lambert Law

• A = εbc
  • A = Absorbance, ε = Molar absorptivity, b = Path length, c = Concentration.
  • Use calibration curves to determine unknown concentrations.

This concludes the Unit 3 review. Stay tuned for Unit 4!