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.
  • $\Delta H = \Sigma H<em>{\text{bonds broken}} - \Sigma H</em>{\text{bonds formed}}$
  • (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: $M<em>1V</em>1 = M<em>2V</em>2$
• Solubility: Like dissolves like.
• Beer-Lambert Law:
  • $A = \epsilon \cdot l \cdot c$
  • A: Absorbance, l: length, c: concentration
• Partial Pressure: $P<em>{\text{total}} = P</em>A + P<em>B + … = n</em>{\text{total}} \frac{RT}{V}$

Chemical Reactions

• General form: $NaCl<em>{(aq)} + AgNO</em>{3(aq)} \rightarrow NaNO<em>{3(aq)} + AgCl</em>{(s)}$
• Limiting Reactant
• OIL RIG (Oxidation Is Loss, Reduction Is Gain)
• Spectator Ions
  • Example: $Cl^-<em>{(aq)} + Ag^+</em>{(aq)} \rightarrow AgCl_{(s)}$
• 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:
  • $\Delta H<em>{\text{vap}}$ (vaporization), $\Delta H</em>{\text{cond}}$ (condensation), $\Delta H<em>{\text{fus}}$ (fusion = melting), $\Delta H</em>{\text{freezing}}$ in kJ/mol - calculate mass.
  • Hess's Law: $\Delta H<em>{\text{rxn}} = \Sigma \Delta H</em>{\text{products}} - \Sigma \Delta H_{\text{reactants}}$
  • Bond Energy: How much energy it takes to break a bond.
    • $\Delta H<em>{\text{rxn}} = \Sigma BE</em>{\text{broken}} - \Sigma BE_{\text{formed}}$
    • 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 $H_2O$.
• $K<em>aK</em>b = K_w$
• Don't need ICE table with acids (use concentration of strong acid).
• $pK<em>a = -\log(K</em>a)$
• $pH = pK_a$ @ 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]