AP Chemistry Speed Review

Introduction

• Presenter: Jeremy Krug
• Purpose: Quick review of major topics in AP Chemistry.
• Resource: Ultimate Review Packet for detailed guides, longer reviews, full exam.
• Price: $24.99 (40% discount for class purchases).

Unit 1: Atoms

• Mole: Unit to count atoms/molecules. 1 mole = atomic/molecular mass in grams.
  • Example: 1 mole of iron = 55.85g, 1 mole of water = 18.02g.
  • Avogadro's number: 6.022 × 10²³ particles per mole.
• Electron Configurations: Stability with 8 electrons in valence shell.
  • Example: Neon = 1s² 2s² 2p⁶.
• Coulomb’s Law: Attraction strength increases with charge and proximity.
• Photoelectron Spectroscopy: Peaks represent sublevels; energy needed to remove electrons.
• Periodic Table Patterns:
  • Atomic radius increases downwards/leftwards.
  • First ionization energy highest at top/right.

Unit 2: Chemical Compounds

• Bonds:
  • Ionic: Metals + nonmetals, electrostatic forces.
  • Covalent: Nonmetals, shared electrons, can be polar/nonpolar.
  • Metallic: Metals, free-moving electrons.
• Structures:
  • Molecules: Covalent, individual units.
  • Ionic Compounds: Lattice, alternating cations/anions.
• Lewis Electron-dot Diagrams: Visualizing molecular shapes and valence electrons.
  • Example Shapes: Tetrahedral (109.5°), Linear (180°), Trigonal Planar (120°).

Unit 3: Intermolecular Forces

• Dispersion Forces: Weak, stronger with larger/more electrons.
• Dipole-Dipole Forces: Polar molecules, stronger than dispersion.
• Hydrogen Bonding: Strong, in molecules with O-H, N-H, F-H.
• States of Matter:
  • Solids: Crystalline, fixed shape/volume.
  • Liquids: Flow, less tightly packed.
  • Gases: Independent molecules, expandable/compressible.
• Gas Laws:
  • Ideal Gas Law: PV=nRT, approximation for small molecules/high temp/low pressure.
  • Temperature Effects: Higher temp = higher kinetic energy.
• Solution Concentration:
  • Molarity: Moles of solute/liters of solution.
  • "Like Dissolves Like" Rule: Polar dissolves polar, nonpolar dissolves nonpolar.

Unit 4: Chemical Reactions

• Equations:
  • Omit spectator ions for net ionic equation.
  • Balance using coefficients, creating mole ratios.
• Reaction Types:
  • Precipitation: Form solid.
  • Oxidation-Reduction: Electron transfer.
  • Acid-Base: Proton transfer.

Unit 5: Kinetics

• Reaction Rates: Relative to balanced equation coefficients.
• Rate Law: Experimentally determined, includes rate constant and reactant orders.
  • Zero, first, and second order relationships.
• Integrated Rate Law: Calculate remaining concentration over time.
• Reaction Mechanism: Slow step determines rate.
• Collision Theory: Molecules must collide with sufficient energy/orientation.
• Catalysts: Lower activation energy.

Unit 6: Thermodynamics

• Heat Transfer Equation: Q = MCΔT.
• Enthalpy (ΔH):
  • Bond enthalpies, formation enthalpies, Hess’s Law.
• Endothermic vs. Exothermic: Heat absorption vs. release.

Unit 7: Equilibrium

• Equilibrium Concept: Forward/reverse rates equal.
• Reaction Quotient (Q) and Equilibrium Constant (K): Q=K at equilibrium.
• ICE Box Method: Organize data for final concentrations or pressures.
• Le Chatelier’s Principle: Reaction response to component changes.

Unit 8: Acids and Bases

• Equations:
  • pH = -log[H⁺], pOH = -log[OH⁻].
  • pH + pOH = 14 at 25°C.
• Strong vs. Weak Acids/Bases: Complete vs. partial ionization.
• Titration: Determine concentration with indicator.
• Buffers: Resist pH changes, calculated with Henderson-Hasselbalch equation.

Unit 9: Applications of Thermodynamics

• Entropy (S): Disorder level, increases from solid to gas.
• Gibbs Free Energy (ΔG): Favorability of a reaction, ΔG = ΔH - TΔS.
• Electrochemistry:
  • Galvanic cells: Anode (oxidation), cathode (reduction).
  • Voltage calculations and Nernst Equation.
  • Electrolysis: Current determines metal plating.

Conclusion

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