AP® Chemistry Cheat Sheet

This comprehensive AP Chemistry Cheatsheet provides essential formulas, key concepts, and critical information across all units of the AP Chemistry curriculum. By summarizing complex topics into easy-to-understand points, it helps students efficiently review and reinforce their knowledge. The inclusion of specific formulas and example problems aids in the application of theoretical concepts, ensuring students are well-prepared for both multiple-choice and free-response questions on the exam. With clear, concise explanations and organized sections, this cheatsheet is an invaluable study aid for achieving a high score on the AP Chemistry exam.

Unit 1: Atomic Structure & Properties

• The Atom: Consists of protons (+), neutrons (0), electrons (-).
• Mole Concept: Relates the mass of an element to the number of particles. One mole equals Avogadro’s number (6.022 x 10²³) particles.
• Molar Mass: The mass of one mole of a substance in grams.
• Isotopes: Atoms of the same element with different numbers of neutrons, leading to different mass numbers.
• Mass Spectroscopy: Identifies the composition of a sample by measuring the mass-to-charge ratio of ions.
• Empirical Formula: The simplest whole number ratio of elements in a compound.
• Electron Configuration: Describes the arrangement of electrons in an atom’s orbitals. Electrons fill orbitals from lowest to highest energy.
• Photoelectron Spectroscopy (PES): Measures the ionization energies of electrons to deduce electronic structure.
• Periodic Trends: Atomic radius, ionization energy, and electronegativity trends across periods and down groups.
• Quantum Mechanical Model: Describes electron distribution in atoms. Orbitals (s, p, d, f) with specific shapes and energy levels.
• Heisenberg Uncertainty Principle: Impossible to know both the position and momentum of an electron simultaneously.
• Example Calculation: Calculate the average atomic mass of an element given the isotopic masses and their abundances.

Unit 2: Molecular & Ionic Bonding

• Intramolecular Forces: Forces within a molecule (ionic, covalent, and metallic bonds).
  • Ionic Bond: Transfer of electrons from a metal to a nonmetal.
  • Polar Covalent Bond: Unequal sharing of electrons between atoms.
  • Nonpolar Covalent Bond: Equal sharing of electrons.
• Ionic Solids: Lattices of cations and anions held together by electrostatic forces.
• Metallic Bonds: Delocalized electrons shared among a lattice of metal atoms.
• Lewis Diagrams: Represent valence electrons and bonds.
• Resonance Structures: Different valid Lewis structures for the same molecule.
• VSEPR Theory: Predicts molecular geometry based on electron pair repulsion.
• Bond Energy: Energy required to break a bond in a molecule.
• Example: Calculate the formal charge of atoms in a molecule to predict the most stable Lewis structure.
• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals (sp, sp², sp³).

Unit 3: Intermolecular Forces & Properties

• Intermolecular Forces: Forces between molecules.
  • London Dispersion Forces (LDFs): Weakest, present in all molecules.
  • Dipole-Dipole Interactions: Between polar molecules.
  • Hydrogen Bonding: Strongest, occurs when H is bonded to F, O, or N.
  • Ion-Dipole Interactions: Between ions and polar molecules.
• Solids:
  • Amorphous Solids: No long-range order (e.g., glass).
  • Crystalline Solids: Ordered structures (ionic, metallic, covalent network, molecular).
• Liquids:
  • Properties: Surface tension, viscosity, capillary action.
  • Laws: Ideal gas law (PV=nRT), law of partial pressures.
• Solutions: Factors affecting solubility (like dissolves like, temperature, pressure).
• Phase Diagrams: Graphs showing conditions (temperature and pressure) at which distinct phases occur and coexist.
• Example: Explain how intermolecular forces affect boiling and melting points of substances.

Unit 4: Chemical Reactions

• Physical vs. Chemical Changes: Physical changes do not alter chemical composition, chemical changes do.
• Precipitation Reactions: Formation of an insoluble product from soluble reactants.
• Net Ionic Equations: Show only the species that undergo a change.
• Balancing Equations: Conservation of mass and charge.
• Stoichiometry: Calculations based on balanced chemical equations.
• Acid-Base Reactions: Transfer of protons between reactants.
  • Titration: Determines the concentration of an unknown acid/base.
• Redox Reactions: Transfer of electrons between reactants.
• Types of Reactions: Synthesis, decomposition, single replacement, double replacement, and combustion reactions.
• Example Problem: Balance a combustion reaction and calculate the amount of product formed.

Unit 5: Kinetics

• Rate of Reaction: Change in concentration of reactants/products over time.
• Rate Laws: Express reaction rate as a function of reactant concentrations.
  • Rate Law Formula: Rate=𝑘[𝐴]𝑚[𝐵]𝑛Rate=k[A]m[B]n
• Reaction Mechanisms: Step-by-step sequence of elementary reactions.
• Collision Model: Reactions occur when particles collide with sufficient energy and correct orientation.
• Catalysts: Lower the activation energy and speed up reactions without being consumed.
• Factors Affecting Reaction Rate: Concentration, temperature, surface area, and catalysts.
• Arrhenius Equation: 𝑘=𝐴𝑒−𝐸𝑎/𝑅𝑇k=Ae−Ea/RT, relates temperature and rate constant.
• Example Calculation: Determine the activation energy given the rate constants at different temperatures.

Unit 6: Thermodynamics

• Kinetic vs. Potential Energy: Kinetic energy relates to motion, potential energy to position or composition.
• Enthalpy (ΔH): Heat absorbed or released in a reaction.
  • Endothermic: Absorbs heat (ΔH > 0).
  • Exothermic: Releases heat (ΔH < 0).
• Calorimetry: Measures heat changes in reactions.
• Hess’s Law: Total enthalpy change is the sum of the enthalpy changes for individual steps.
• Standard Enthalpy of Formation (ΔH°f): Enthalpy change when one mole of a compound forms from its elements.
• Second Law of Thermodynamics: Entropy of an isolated system always increases.
• Example: Use Hess’s Law to calculate the enthalpy change of a reaction from known enthalpy changes of related reactions.

Unit 7: Equilibrium

• Dynamic Equilibrium: Rate of the forward reaction equals the rate of the reverse reaction.
• Equilibrium Constant (K): Ratio of product concentrations to reactant concentrations at equilibrium.
• Le Chatelier’s Principle: A system at equilibrium will adjust to relieve any applied stress.
• ICE Tables: Used to calculate changes in concentrations of reactants/products.
• Le Chatelier’s Principle Examples: Predict the shift in equilibrium when concentration, temperature, or pressure changes.
• Equilibrium Expressions: Write Kc and Kp expressions for given reactions.

Unit 8: Acids & Bases

• Acid-Base Definitions: Bronsted-Lowry (proton transfer), Lewis (electron pair).
• pH and pOH: Measures of acidity/basicity.
  • pH = -log[H₃O⁺]
  • pOH = -log[OH⁻]
• Titrations: Determine the concentration of an acid/base by neutralization.
• Buffers: Solutions that resist changes in pH when small amounts of acid or base are added.
• Strength of Acids and Bases: Strong acids/bases dissociate completely, weak acids/bases do not.
• pKa and pKb: Measures of acid and base strength.
• Example Calculation: Calculate the pH of a weak acid solution using the acid dissociation constant (Ka).

Unit 9: Applications of Thermodynamics

• Entropy (S): Measure of disorder.
• Gibbs Free Energy (ΔG): Determines spontaneity of a reaction.
  • ΔG = ΔH – TΔS
  • Spontaneous when ΔG < 0.
• Electrochemistry: Study of redox reactions and their applications.
  • Galvanic Cells: Convert chemical energy to electrical energy.
  • Electrolytic Cells: Use electrical energy to drive non-spontaneous reactions.
• Standard Electrode Potentials: Determine cell potential from standard reduction potentials.

These additions provide more depth and context, making the cheat sheet even more helpful for students preparing for the AP Chemistry exam.