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(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.