Overview
This lecture provides a rapid review of the essential content and problem-solving strategies needed for the AP Chemistry exam, covering key concepts from all course units.
Unit 1: Composition of Compounds & Atomic Structure
- Mass percent of an element in a compound = (mass of element in formula / total molar mass) × 100.
- Patterns: lighter cation → higher anion mass percent; heavier cation → higher cation mass percent.
- Mass spectrometry: bar heights show isotope abundance; average atomic mass calculated from weighted mean.
- Electron configuration: add/remove electrons for ions; cations lose from outermost shell, anions gain to outer shell.
- Atomic radius: increases down and left due to more shells and fewer protons.
- Ionic radius: cations smaller, anions larger; more protons in isoelectronic series means smaller radius.
- First ionization energy: increases up and right due to less shielding and greater nuclear charge.
- Photoelectron spectroscopy (PES): peak position = sublevel, height = # electrons; more protons → greater binding energy.
Unit 2: Bonding & Molecular Structure
- Lewis diagrams: central atom in the middle; fill octets using shared pairs and lone pairs.
- Sigma bonds: single = 1 sigma; double = 1 sigma + 1 pi; triple = 1 sigma + 2 pi.
- Hybridization: count electron regions; 4 = sp3, 3 = sp2.
- Molecular geometry: 4 regions = tetrahedral (109.5°); 3 = trigonal planar (120°); 2 + 2 lone pairs = bent (104.5°).
- Polarity: molecule is polar if there's an unbalanced negative region; always has London forces, polar = dipole-dipole (+ H-bonding if H–F/O/N).
- Melting point: higher ion charge and smaller ions increase melting point (Coulomb’s law).
- Ionic solids dissolve in water via ion-dipole interactions if those forces exceed lattice energy.
Unit 3: Gases and Solutions
- Ideal gas law: PV = nRT; solve for unknowns using given values.
- For mixtures: total moles = sum of partial moles; molar mass = grams/moles.
- Spectrophotometry: use Beer-Lambert Law; calibration curve links absorbance to concentration; outliers suggest contamination.
Unit 4: Reactions and Stoichiometry
- Net ionic equations: include only ions involved in the reaction; omit spectator ions.
- Redox: oxidation = lose electrons; reduction = gain electrons; balance electrons in half-reactions.
- Balancing reactions: ensure atoms and charge conservation.
- Precipitation: insoluble product forms solid; write only reacting ions in net ionic.
- Stoichiometry steps: convert to moles, use mole ratio, convert to final units.
Unit 5: Kinetics
- Determine reaction order by comparing trials where only one reactant changes.
- Rate law: rate = k[A]^m[B]^n; overall order = sum of exponents.
- Rate constant units depend on overall order (e.g., M⁻²s⁻¹ for third order).
- Graphing: [A] vs. time (zero order), ln[A] vs. time (first order), 1/[A] vs. time (second order); straight line identifies order.
- Reaction mechanism: intermediate is made then used; rate law from slow (rate-determining) step.
Unit 6: Thermochemistry
- Heat transfer: Q = mc∆T; negative Q = heat lost, positive Q = gained.
- Heating/cooling curves: temperature constant during phase changes.
- ∆H (enthalpy change) = ΣHf(products) − ΣHf(reactants); solve for missing values algebraically.
Unit 7: Equilibrium
- K expressions: products/reactants (omit solids/liquids); Kc uses concentrations, Kp uses pressures.
- ICE tables help track initial, change, equilibrium values.
- K << 1: reactants favored; K >> 1: products favored.
- Q vs. K: Q < K, reaction shifts right; Q > K, reaction shifts left.
- Le Chatelier’s Principle: system shifts to counter changes (add/remove, pressure, temp).
- Endothermic: heat as reactant; exothermic: heat as product.
Unit 8: Acids and Bases
- [H₃O⁺][OH⁻] = 1×10⁻¹⁴ at 25°C; pH = -log[H₃O⁺], pOH = -log[OH⁻]; pH + pOH = 14.
- Strong acids: concentration = [H₃O⁺]; strong bases: Group 1/2 hydroxides, multiply [OH⁻] accordingly.
- Weak acids/bases: use ICE table and Ka/Kb; neglect x in denominator for small x.
- Percent dissociation = (x / initial) × 100.
- Acid/base nature of salt: compare parent acid/base strengths using ion swap trick.
- Titration: MaVa = MbVb at equivalence; at half-equivalence, pH = pKa.
Unit 9: Thermodynamics and Electrochemistry
- Entropy (S): solid < liquid < aqueous < gas; more molecules or higher T = more S.
- Thermodynamic favorability: exothermic (−∆H) and increased entropy (+∆S) always favored.
- ∆G = ∆H – T∆S; negative ∆G = favored.
- Galvanic/voltaic cells: E°cell = E°cathode − E°anode; positive E°cell = favored.
- Electrons flow anode to cathode; cations → cathode, anions → anode via salt bridge.
Key Terms & Definitions
- Mass percent — percent by mass of an element in a compound.
- Ionization energy — energy needed to remove an outer electron.
- Isoelectronic — species with same number of electrons.
- Sigma bond — single covalent bond.
- Hybridization — mixing of atomic orbitals (sp, sp2, sp3).
- Polarity — separation of electric charge in a molecule.
- Spectrophotometry — measuring light absorbance to find concentration.
- Stoichiometry — quantitative relationships in chemical reactions.
- Intermediate — produced then consumed in reaction mechanism.
- Enthalpy (∆H) — heat change at constant pressure.
- Entropy (S) — measure of disorder or dispersal.
- Gibbs free energy (∆G) — determines favorability, ∆G < 0 means favored.
- Galvanic cell — electrochemical cell producing electricity from a redox reaction.
Action Items / Next Steps
- Review and practice problems using supplied guided notes PDF.
- Focus study on weak vs. strong acid/base problems, equilibrium calculations, and energy diagrams.
- Prepare calculator and reference sheets for exam day.