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CHEM 105 Exam 2 Printable Flashcard

Optimized two-sided reference card featuring a 2-column VSEPR table (all 13 shapes SN 2–6 fully visible), lone-pair placement rules, Lewis structures & octet exceptions, formal charge rules, bond polarity (ΔEN), and Molecular Orbital (MO) Theory. Headers and footers removed to maximize usable reference space.

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VSEPR: Steric Numbers 2, 3 & 4 (6 Shapes) LP: LP-LP > LP-BP > BP-BP
SN B/LP Electron Domain Molecular Shape Angles
2 2 / 0 Linear Linear 180°
3 3 / 0 Trigonal Planar Trig. Planar 120°
2 / 1 Trigonal Planar Bent (V-shape) <120° (118°)
4 4 / 0 Tetrahedral Tetrahedral 109.5°
3 / 1 Tetrahedral Trig. Pyramidal <109.5° (107°)
2 / 2 Tetrahedral Bent (Angular) <109.5° (104.5°)
VSEPR: Steric Numbers 5 & 6 (7 Shapes) Angles & Shapes
SN B/LP Electron Domain Molecular Shape Angles
5 5 / 0 Trig. Bipyramid Trig. Bipyramidal 90°, 120°
4 / 1 Trig. Bipyramid Seesaw <90°, <120°
3 / 2 Trig. Bipyramid T-shaped <90° (87.5°)
2 / 3 Trig. Bipyramid Linear 180°
6 6 / 0 Octahedral Octahedral 90°, 180°
5 / 1 Octahedral Square Pyramidal <90° (85°)
4 / 2 Octahedral Square Planar 90°, 180°
LP Repulsion & Angle Compression:
• Repulsion hierarchy: LP–LP > LP–BP > BP–BP.
• Lone pairs take up more space, compressing adjacent bond angles:
  CH₄ (0 LP) = 109.5° → NH₃ (1 LP) = 107.3° → H₂O (2 LP) = 104.5°.
Lone Pair Placement Rules:
• SN = 5 (Trig. Bipyramidal): Lone pairs always occupy equatorial positions (120° apart) to minimize 90° repulsions. Never axial!
• SN = 6 (Octahedral): 1st LP anywhere; 2nd LP always goes opposite (axial, 180°), forming Square Planar.
Lewis, Formal Charge & ΔEN FC = V − L − B
  • 1. Valence e⁻ Sum: Add all valence e⁻ (anions: add charge; cations: subtract).
  • 2. Central Atom: Least EN (never H; halogens terminal unless oxoacid).
  • 3. Duets & Octets: Single bonds (2e⁻) first; complete outer duets (H=2) & octets; place leftovers on central; form multiple bonds if central lacks octet.
  • 4. Octet Exceptions: Incomplete: H(2), Be(4 in BeCl₂), B(6 in BF₃). Expanded: Period 3+ (P, S, Cl, Br, I, Xe) use d-orbitals to hold 10–12e⁻ (PCl₅, SF₆, XeF₄).
  • 5. Formal Charge: FC = V − L − B (Valence − Dots − Sticks). ∑FC = net charge.
  • 6. Best Structure: (1) FC closest to 0; (2) −FC on most EN atom (F>O>N>Cl); (3) No like FC (+/+ or −/−) on adjacent bonded atoms.
  • 7. Bond Polarity (ΔEN):
    • ΔEN < 0.4: Nonpolar Covalent (equal share, no dipole; C−H, C−C).
    • 0.4 ≤ ΔEN < 1.8: Polar Covalent (unequal share, δ⁺ → δ⁻, μ>0; H−Cl, C−O).
    • ΔEN ≥ 1.8: Ionic (electron transfer, cation/anion lattice; NaCl, CaO).
Molecular Orbital (MO) Theory BO = ½(Nb − Na)
  • • Bond Order: BO = ½(N_b − N_a) = ½(Bonding e⁻ − Antibonding e⁻).
    • BO > 0 → Stable. Higher BO → stronger, shorter bond. BO = 0 → unstable (He₂).
  • • Magnetism:
    • Paramagnetic: ≥1 unpaired e⁻ → attracted to magnetic field (O₂, B₂).
    • Diamagnetic: All e⁻ paired → weakly repelled by magnetic field (N₂, C₂, F₂, CO).
  • • Diatomic MO Diagrams & Filling Order:
    1. O₂, F₂, Ne₂ (Normal, no sp-mixing):
      σ₂s < σ*₂s < σ₂p < (π₂p=π₂p) < (π*₂p=π*₂p) < σ*₂p
      O₂ (12 val e⁻): 2 unpaired e⁻ in π*₂p (Hund's rule) → Paramagnetic! BO = 2.
    2. B₂, C₂, N₂ (sp-Mixing Inversion):
      σ₂s < σ*₂s < (π₂p=π₂p) < σ₂p < (π*₂p=π*₂p) < σ*₂p
      π₂p drops lower than σ₂p due to 2s-2p orbital mixing!
    3. CO (Heteronuclear): 10 val e⁻; N₂-like diagram → BO = 3, Diamagnetic.
  • • AOs → MOs: 2 AOs combine → 2 MOs: 1 bonding (lower energy), 1 antibonding (*) with nodal plane (zero e⁻ probability).