AP Chemistry Speed Review by Jeremy Krug

Introduction

• A quick review of major AP Chemistry topics in under 20 minutes.
• Ultimate Review Packet: comprehensive study guides, longer videos, full-length exam.
• Visit UltimateReviewPacket.com for more details.

Unit 1: Atoms

• Mole Concept: One mole = atomic mass (or sum of atomic masses for compounds) in grams.
  • Example: 1 mole of iron = 55.85 g, 1 mole of H2O = 18.02 g.
  • Avogadro's Number: 6.022 x 10^23 particles.
• Electron Configurations: Neon is 1s² 2s² 2p⁶.
• Stability: Atoms are stable with 8 electrons in the valence shell (octet rule).
• Coulomb's Law:
  • Greater charges = stronger attraction.
  • Closer distance = stronger attraction.
• Photoelectron Spectroscopy: Represents energy to remove electrons.
• Periodic Table Patterns:
  • Atomic radius increases downwards and left.
  • Ionization energy is highest at top right.
  • Anions increase in size, cations decrease.

Unit 2: Chemical Compounds

• Ionic Bonds: Electrostatic attraction between metals and nonmetals.
• Covalent Bonds: Electrons shared between nonmetals.
  • Polar (unequal sharing) vs. Nonpolar (equal sharing).
• Metallic Bonds: Free-flowing electrons in metal ions.
• Lewis Structures: Visualize molecules; aim for octet, may form double/triple bonds.
• Molecular Geometry:
  • Tetrahedral (109.5°), Linear (180°), Trigonal Planar (120°).

Unit 3: Intermolecular Forces

• Dispersion Forces: Weak, increase with size and electrons.
• Dipole-Dipole: Attraction between polar molecules.
• Hydrogen Bonding: Strong forces in molecules with O-H, N-H, or F-H bonds.
• States of Matter:
  • Solids (crystalline), Liquids (flow), Gases (expand/compress).
• Ideal Gas Law: PV=nRT, describes gas behavior.
• Kinetic Energy: Increases with temperature.
• Solution Concentration: Molarity (moles solute/liters solution).
• Spectrophotometry: Analyzes solution concentration via absorbance.

Unit 4: Chemical Reactions

• Net Ionic Equations: Omit spectator ions.
• Balancing Equations: Ensure equal atoms on both sides.
• Stoichiometry:
  • Convert quantities to moles.
  • Use mole ratio for conversions.
  • Types: Precipitation, Redox, Acid-Base.

Unit 5: Kinetics

• Rate Laws: Determine reaction rates experimentally.
  • Zero, First, Second order reactions.
• Integrated Rate Law: Calculate reactant amounts over time.
• Reaction Mechanism: Series of steps, slow step determines rate.
• Temperature/Collision Theory:
  • Increase temperature, decrease particle size, increase concentration, use catalysts.

Unit 6: Thermodynamics

• Heat Transfer: Q = MCΔT (Joules, mass, specific heat, temperature change).
• Enthalpy (ΔH): Reaction heat change (kJ/mole).
  • Bond enthalpies, enthalpy of formation, Hess's Law.
• Exothermic vs. Endothermic: Releases vs. absorbs heat.

Unit 7: Equilibrium

• Dynamic Equilibrium: Forward/reverse rates equal.
• Reaction Quotient (Q): Products/reactants concentration ratio.
• Equilibrium Constant (K): High K = more product; low K = more reactant.
• Le Chatelier’s Principle: Reaction shifts with component changes.
• ICE Box Method: Solve equilibrium problems.

Unit 8: Acids and Bases

• pH and pOH: -log of ion concentrations.
• Kw: Ion product of water at 25°C.
• Strong vs. Weak Acids/Bases: Degree of ionization.
• Titrations: Determine concentration via endpoint using indicators.
• Buffers: Resist pH changes; use Henderson-Hasselbalch equation.

Unit 9: Applications of Thermodynamics

• Entropy (S): Measure of disorder; solids < liquids < gases.
• Gibbs Free Energy (ΔG): Favorability measure; ΔG = ΔH - TΔS
• Electrochemistry:
  • Galvanic cells: Oxidation at anode, reduction at cathode.
  • Nernst Equation for non-standard conditions.
  • Electrolysis: External source drives reactions.

Conclusion

• Comprehensive review of AP Chemistry foundational topics.
• Encourage use of additional resources for thorough understanding.