AP Chemistry Speed Review

Introduction

• Presenter: Jeremy Krug
• Purpose: Quick review of major AP Chemistry topics in less than 20 minutes.
• Additional resources available at UltimateReviewPacket.com, including study guides, review videos, and a full-length exam.

Unit 1: Atoms

• Mole concept: Used to count particles, equals atomic mass in grams (e.g., 1 mole of iron = 55.85g).
• Electron configurations: Example, Neon: 1s2, 2s2, 2p6.
• Stability: Atoms are stable with 8 electrons (octet) in the valence shell.
• Coulomb’s Law: Attraction between opposite charges increases with greater charge and proximity.
• Photoelectron spectroscopy: Peaks represent sublevels; taller peaks have more electrons.
• Periodic Table Patterns:
  • Atomic radius increases downwards and leftwards.
  • First ionization energy highest at the top-right.
  • Anions are larger; cations are smaller.

Unit 2: Chemical Compounds

• Ionic bonds: Form between metals and nonmetals; electrostatic attraction.
• Covalent bonds: Form between nonmetals; can be polar (unequal sharing) or nonpolar (equal sharing).
• Metallic bonding: Positive metal ions in a sea of electrons.
• Lewis diagrams: Visualize molecule shapes; aim for eight valence electrons.
• Molecular geometry:
  • Tetrahedral: 109.5°
  • Linear: 180°
  • Trigonal planar: 120°

Unit 3: Intermolecular Forces

• Dispersion forces: Weak, increase with size and electrons.
• Dipole-dipole forces: Stronger than dispersion, occur between polar molecules.
• Hydrogen bonding: Strong, involves O-H, N-H, or F-H bonds.
• States of Matter:
  • Solids: Crystalline, fixed shape/volume.
  • Liquids: Molecules slide; flows.
  • Gases: Independent molecules, compressible.
• Ideal Gas Law: PV=nRT; describes gas relationships.
• Solution Concentration: Molarity = moles/volume.
• Spectrophotometry: Measures concentration via light absorption.

Unit 4: Chemical Reactions

• Net ionic equations: Exclude spectator ions.
• Balancing equations: Ensure equal atoms on both sides.
• Reaction types:
  • Precipitation: Formation of a solid.
  • Redox: Electron transfer, oxidation and reduction.
  • Acid-base: Proton exchange, forming conjugates.

Unit 5: Kinetics

• Rate laws: Described by reactant concentration and reaction orders.
• Integrated rate laws: Calculate remaining reactant over time.
• Reaction mechanisms: Multiple steps; slowest step determines rate.
• Catalysts: Lower activation energy, speed up reactions.

Unit 6: Thermodynamics

• Heat transfer: Q = M C ΔT.
• Enthalpy (ΔH): Heat change in reactions; calculated via bond enthalpies or Hess’s Law.
• Exothermic vs Endothermic: Heat release vs absorption.

Unit 7: Equilibrium

• Equilibrium state: Rates of forward and reverse reactions equal.
• Le Chatelier’s Principle: Reactions shift to counteract changes.
• Equilibrium constant (K): Product/reactant concentration ratio.

Unit 8: Acids and Bases

• pH and pOH: Related to concentration of H+ and OH- ions.
• Strong vs weak acids/bases: Complete vs partial ionization.
• Titration: Determines concentration via endpoint.
• Buffers: Resist pH changes; calculated by Henderson-Hasselbalch equation.

Unit 9: Applications of Thermodynamics

• Entropy (S): Measure of disorder; gases > solutions > liquids > solids.
• Gibbs Free Energy (ΔG): Indicates reaction favorability.
• Electrochemistry: Galvanic cells involve oxidation/reduction reactions.
• Nernst Equation: Calculates cell voltage under non-standard conditions.

Conclusion

• Major points of the entire AP Chemistry course summarized.
• Encouragement to use additional resources for thorough review.