AP Chemistry Notes

Atoms and Molecular Geometry

  • Basic Geometries
    • Tetrahedral: 4 bonding pairs, 0 lone pairs, bond angle 109.5°
    • Trigonal Pyramidal: 3 bonding pairs, 1 lone pair, bond angle ~109.5°
    • Bent: 2 bonding pairs, 2 lone pairs, bond angle ~109.5°
    • Trigonal Planar: 3 bonding pairs, 0 lone pairs, bond angle 120°
    • Bent/Angular: 2 bonding pairs, 1 lone pair, bond angle ~120°
    • Linear: 2 bonding pairs, 0 lone pairs, bond angle 180°
    • Trigonal Bipyramidal: 5 bonding pairs, 0 lone pairs, bond angles 90° & 120°
    • See-saw: 4 bonding pairs, 1 lone pair, bond angles ~90° & 120°
    • T-shaped: 3 bonding pairs, 2 lone pairs
    • Octahedral: 6 bonding pairs, 0 lone pairs, bond angle 90°
    • Square Pyramidal: 5 bonding pairs, 1 lone pair, bond angle ~90°
    • Square Planar: 4 bonding pairs, 2 lone pairs, bond angle 90°

Intermolecular Forces and Radiation

  • Hydrogen Bonding: Occurs in molecules containing O-H, N-H, or F-H bonds.
  • Microwave Radiation: Causes rotation of molecules.
  • Infrared Radiation: Causes vibration of molecules.
  • UV / Visible Radiation: Causes electrons to transition to different energy levels.

Solubility Rules

  • Compounds containing Na^+, K^+, or NH_4^+ ions are always soluble in water.
  • Compounds containing NO_3^- ions are always soluble in water.

Bond Enthalpies

  • The change in enthalpy (\Delta H) is calculated as the sum of the enthalpies of bonds broken minus the sum of the enthalpies of bonds formed.
    • \Delta H = \Sigma H{\text{bonds broken}} - \Sigma H{\text{bonds formed}}
    • (LEFT SIDE - RIGHT SIDE)

Atomic Structure and Properties

  • Electronegativity: The force attracting electrons (F highest).
  • Ionization Energy: The energy required to lose an electron (F highest).
  • Electron Affinity: The energy change after adding an electron (F-).

Bonding

  • Intramolecular Bonding
    • Ionic: Between metal & non-metal (e.g., Na^+Cl^-, Na^+ Cl^-).
    • Covalent: Between two non-metals.
    • Metallic: Between two metals (alloys). Electrons are not associated with a single atom but are "delocalized electrons" forming a "sea of electrons".
      • Substitutional: Atoms of compatible radii.
      • Interstitial: Atoms of different radii where smaller atoms fill spaces.
    • Polar Covalent: Difference in electronegativity forms partial charges.

Intermolecular Forces and Properties

  • London Dispersion Forces (LDF): Induced dipole forces due to temporary imbalances.
  • Polarizability: How easy it is to polarize a molecule. Higher polarizability with larger electron cloud.
    • Stronger LDF leads to higher boiling point, longer molecules/surface area, or larger molar mass.
  • Dipole-Dipole: Permanent dipole.
  • Hydrogen Bonds: H bonded to F, O, or N - Strongest IMFs - Highest B.P.
  • Ion-Dipole

Vapor Pressure

  • Lower boiling point results in higher vapor pressure.

Gas Laws

  • Ideal Gas Law: (not specified in provided text)
  • Molarity equation: M1V1 = M2V2
  • Solubility: Like dissolves like.
  • Beer-Lambert Law:
    • A = \epsilon \cdot l \cdot c
    • A: Absorbance, l: length, c: concentration
  • Partial Pressure: P{\text{total}} = PA + PB + … = n{\text{total}} \frac{RT}{V}

Chemical Reactions

  • General form: NaCl{(aq)} + AgNO{3(aq)} \rightarrow NaNO{3(aq)} + AgCl{(s)}
  • Limiting Reactant
  • OIL RIG (Oxidation Is Loss, Reduction Is Gain)
  • Spectator Ions
    • Example: Cl^-{(aq)} + Ag^+{(aq)} \rightarrow AgCl_{(s)}
  • Brønsted-Lowry: Acid = H^+ donor (proton), Base = proton / H^+ acceptor.

Kinetics

  • Rate Law: Rate = k[A]^x[B]^y
  • Half-life follows first-order rate law.
  • Use coefficients as orders only in elementary steps.
  • Intermediate - created then used up.
  • The rate-determining step is the slow step.
    • Don't use intermediates in the rate law.

Thermodynamics

  • 1 \text{ Cal} = 1 \text{ kcal} = 1000 \text{ calories}
  • Calorie (food)= Cal
  • Phase Changes:
    • \Delta H{\text{vap}} (vaporization), \Delta H{\text{cond}} (condensation), \Delta H{\text{fus}} (fusion = melting), \Delta H{\text{freezing}} in kJ/mol - calculate mass.
    • Hess's Law: \Delta H{\text{rxn}} = \Sigma \Delta H{\text{products}} - \Sigma \Delta H_{\text{reactants}}
    • Bond Energy: How much energy it takes to break a bond.
      • \Delta H{\text{rxn}} = \Sigma BE{\text{broken}} - \Sigma BE_{\text{formed}}
      • Divide by mol of rxn if needed.

Equilibrium

  • K_{eq} only changes with temperature.
  • Q can be calculated at any moment to find direction.
  • Le Châtelier's Principle: (not specified in provided text)
    • If pressure moves away from the most gaseous side, the system equilibrates.
    • If Q > K_{eq}, the reaction is oversaturated.
      • Equilibrium shifts in the opposite direction.
  • Common Ion Effect: Adding concentration of a common ion.

Acids and Bases

  • Polyprotic acids - focus on the first equation.
  • Strong acids and bases dissolve in H_2O.
  • KaKb = K_w
  • Don't need ICE table with acids (use concentration of strong acid).
  • pKa = -\log(Ka)
  • pH = pK_a @ half equivalence point = [weak acid] = [conjugate base].
  • Buffer:
    • pH = pK_a + \log \frac{[A^-]}{[HA]}
    • Favorable at low temps if -\Delta H and +\Delta S. Unfavorable +\Delta H and -\Delta S.
    • Buffers: [weak acid] & [conjugate base]