(Jeremy Krug) AP Chem Unit 3 Review | Properties of Substances and Mixtures

Intermolecular Forces

• London Dispersion Forces

  • Present in all molecules, typically the weakest forces.

  • Strength increases with larger molecules due to more electrons.

  • Only intermolecular force in nonpolar molecules.

• Dipole-Dipole Forces

  • Occur between polar molecules, where the positive pole of one molecule attracts the negative pole of another.

  • Stronger than London dispersion forces.

• Hydrogen Bonding

  • A special case of dipole-dipole forces, occurs when hydrogen is bonded to oxygen, nitrogen, or fluorine.

  • Strongest type of intermolecular force.

• Ion-Dipole Forces

  • Occur in solutions where polar molecules (like water) interact with ions in ionic compounds.

  • Can cause ionic compounds to dissolve in polar solvents if the ion-dipole forces are stronger than ionic forces.

Properties of Solids

• Ionic Solids

  • High melting points, brittle, conduct electricity when dissolved in water due to attraction between charged ions.

• Covalent Network Solids

  • Extremely strong due to extensive covalent bonding in all directions (e.g., diamond and silicon dioxide).

• Molecular Solids

  • Made of individual molecules with relatively weak forces, leading to lower melting points (e.g., sugar).

• Metallic Solids

  • Composed of metals or alloys, malleable, ductile, and good conductors of electricity due to a 'sea of electrons'.

• Crystalline Solids

  1. Definition: Have a well-defined, orderly structure.

  2. Characteristics:

     • Regular arrangement of particles.

     • Sharp melting points.

     • Examples: Salt, diamonds, ice

• Amorphous Solids

  1. Definition: Lack a regular arrangement of particles.

  2. Characteristics:

     • No defined melting point, deform gradually when heated.

     • Examples: Glass, plastics, gels.

States of Matter

• Solids: Particles are close together, with limited motion (only vibrational).

• Liquids: Particles are further apart, allowing them to flow and slide past each other.

• Gases: Particles are far apart and move independently, allowing them to expand and be compressed.

Ideal Gas Law

• Formula: PV = nRT

  • P = pressure (atm), V = volume (L), n = number of moles, R = Universal Gas Constant (0.08206 L atm K⁻¹ mol⁻¹), T = temperature (K).

  • Use this equation to solve for unknown variables using provided constants.

Gas Mixtures

• Partial Pressure: The total pressure is the sum of individual gas pressures; calculate the partial pressure using the mole fraction.

Temperature and Kinetic Energy

• Temperature reflects the average kinetic energy of molecules; higher temperature indicates faster-moving molecules.

• Use the Boltzmann distribution to visualize the range of molecular speeds at different temperatures.

Ideal vs. Real Gases

• Ideal gases: No intermolecular attractions, negligible volume.

• Real gases, such as helium, behave ideally under high temperature and low pressure where particle interaction is minimized.

Mixture Types

• Heterogeneous Mixtures: Visible individual components.

• Homogeneous Mixtures (Solutions): Uniformly distributed components; focus on molarity (M = moles of solute/L of solution).

Separating Solutions

• Distillation: Capitalizes on differing boiling points to separate components by boiling and condensing.

• Chromatography: Involves separation based on different adherence rates to a stationary phase, producing faster or slower passage through a mobile phase.

Solubility Principles

• Rule of Thumb: "Like dissolves like"; polar vs. nonpolar solvents and solutes.

Electromagnetic Spectrum and Molecules

• Different portions of the spectrum affect molecules differently:

  • UV/Visible Light: Causes electronic transitions.

  • Infrared Radiation: Causes vibrational energy changes.

  • Microwave Radiation: Induces rotational transitions.

Light and Photons

• Dual Nature of Light: Can behave as waves and particles (photons).

• Photonic Calculations:

  • Use c = λν (speed of light = wavelength x frequency) to relate wavelength and frequency.

  • Use E = hν (energy = Planck’s constant x frequency) to find energy of a photon (h = 6.626 x 10⁻³⁴ J s).

Spectrophotometry and Beer-Lambert Law

• Beer-Lambert Law: A = εbc

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

  • Construct calibration curves to determine unknown concentrations from absorbance measurements.

• Outliers: Possible contamination or dilution can lead to data points straying from the expected curve.