AP Chemistry Speed Review by Jeremy Krug

Introduction

• AP Chemistry overview in less than 20 minutes.
• Not a replacement for a full AP course, but a good starting point for review.
• Ultimate Review Packet ($24.99, 40% class discount available).
  • Includes study guides, longer review videos, and a full-length exam.

Unit 1: Atoms

• The Mole:
  • Used to count large numbers of atoms/molecules.
  • A mole of element = atomic mass in grams.
  • 1 mole = 6.022 x 10²³ particles (Avogadro's number).
  • Example: 1 mole of iron ≈ 55.85g, 1 mole of water ≈ 18.02g.
• Electron Configurations:
  • Neon: 1s² 2s² 2p⁶.
  • Atoms most stable with 8 valence electrons.
  • Coulomb's Law: Attraction force ∝ magnitude of charge, inversely ∝ distance.
• Photoelectron Spectroscopy:
  • Peaks represent sublevels; energy related to electron removal.
• Periodic Table Patterns:
  • Atomic radius increases down and left.
  • First ionization energy highest at top right.
• Ions:
  • Gaining electrons = anions (larger), losing electrons = cations (smaller).

Unit 2: Chemical Compounds

• Ionic Bonds:
  • Between metals and nonmetals; electrostatic forces.
• Covalent Bonds:
  • Between nonmetals; sharing electrons.
  • Polar (unequal) vs. nonpolar (equal).
• Metallic Bonds:
  • Metals/alloys; electrons move freely.
• Lewis Diagrams:
  • Visualize molecule shapes; exceptions exist.
  • Shapes: tetrahedral (109.5°), linear (180°), trigonal planar (120°).

Unit 3: Intermolecular Forces

• Dispersion Forces:
  • Weak, increase with size/electrons.
• Dipole-Dipole Forces:
  • Polar molecules; stronger than dispersion.
• Hydrogen Bonding:
  • Strong forces in O-H, N-H, F-H bonds.
• States of Matter:
  • Solids (fixed shape/volume), liquids (flow), gases (expand/compress).
• Ideal Gas Law (PV=nRT):
  • Describes gas relationships; ideal conditions vary.
• Kinetic Energy and Temperature:
  • Higher temperature = higher average kinetic energy.

Unit 4: Chemical Reactions

• Net Ionic Equations:
  • Omit spectator ions.
• Balancing Equations:
  • Use coefficients for mole ratios.
• Reaction Types:
  • Precipitation, oxidation-reduction, acid-base.

Unit 5: Kinetics

• Rate Laws:
  • Rate = k[Reactant]^(order).
  • Order impacts rate change with concentration.
• Reaction Mechanisms:
  • Multiple steps; slowest determines rate.
• Collisions:
  • Must have sufficient energy/orientation.
• Catalysts:
  • Lower activation energy.

Unit 6: Thermodynamics

• Heat Transfer (Q=MCDeltaT):
  • Q = heat (J), M = mass (g), C = specific heat, DeltaT = temp change.
• Enthalpy (Delta H):
  • Bond enthalpies or formation enthalpies.
• Hess’s Law:
  • Delta H of reactions can be summed.

Unit 7: Equilibrium

• Reaction Quotient (Q):
  • Products/reactants ratio at given point.
• Equilibrium Constant (K):
  • Large K = more product; small K = more reactant.
• Le Chatelier’s Principle:
  • Changes shift equilibrium position.

Unit 8: Acids and Bases

• pH and pOH:
  • pH + pOH = 14 at 25°C.
• Equilibrium Problems:
  • ICE boxes for weak acids/bases.
• Titrations and Buffers:
  • Indicators show endpoints, buffer calculations via Henderson-Hasselbalch.

Unit 9: Applications of Thermodynamics

• Entropy (S):
  • Disorder measure; gases > solutions > liquids > solids.
• Gibbs Free Energy (Delta G):
  • Thermodynamic favorability; related to K.
• Electrochemistry:
  • Galvanic cells, anode/cathode roles, voltage calculations.

Conclusion

• Overview of the entire AP Chemistry course.
• Encouragement to use Ultimate Review Packet for thorough preparation and additional resources.