Bonding
High-Yield Summary
- Atoms bond via valence electrons, generally following the octet rule (8 valence electrons, noble-gas-like configuration).
- Electronegativity difference (ΔEN) classifies bond type: <0.5 nonpolar covalent, 0.5–1.7 polar covalent, >1.7 ionic.
- Ionic bonds form mainly between metals (lose e⁻ → cations) and nonmetals (gain e⁻ → anions); electrons are fully transferred, not shared.
- Ionic compounds pack into an ordered 3D crystal lattice (e.g., NaCl: 6:6 Na⁺/Cl⁻ coordination), giving high melting/boiling points.
- Lattice energy (Coulomb's law, F = kq₁q₂/r²) rises with greater ionic charge and smaller ionic radius — charge has the larger effect.
Octet Rule Exceptions
- Incomplete octet: H, He (stable at 2); Li (2); Be (4); B (6) — Be/B as central atom often gives an electron-deficient molecule (BeCl₂, BCl₃).
- Expanded octet: any period 3+ element (empty d-orbitals) — P can hold 10, S 12, Cl 14.
- Odd-electron species: total valence electrons is odd, so no arrangement gives every atom a full octet (e.g., NO: 5 + 6 = 11).
- Reliable octet-followers: C, N, O, F, Na, Mg.
Electronegativity Difference (ΔEN) → Bond Type
ΔEN < 0.5 → nonpolar covalent | 0.5 ≤ ΔEN ≤ 1.7 → polar covalent | ΔEN > 1.7 → ionic
- ΔEN = Difference in electronegativity between the two bonding atoms (Pauling scale)
- Mnemonic: MeTals lose electrons to become caTions (posiTive); Nonmetals gain electrons to become aNions (Negative).
Coulomb's Law (Lattice Energy)
F = k(q₁q₂) / r²
- q₁, q₂ = Charges of the two ions
- r = Distance between ion centers
- k = Proportionality constant
- Lattice energy ↑ as charge ↑ or radius ↓; charge (numerator, q₁q₂) dominates over radius (denominator, r²).
- Example: MgO (2+, 2−) has a charge product 4× that of NaCl (1+, 1−) → much higher lattice energy → MgO melts at ~2,800°C vs. NaCl's ~801°C.
Common MCAT Trap
- Ionic bonds don't share electrons at all — they're fully transferred. Don't describe an ionic bond as 'unequal sharing' (that's polar covalent).
- Charge, not size, dominates lattice energy — a compound with smaller ions but lower charge can still have weaker lattice energy than one with larger ions but higher charge.
- Boron and beryllium violating the octet rule (electron-deficient) is easy to forget since C/N/O/F next door obey it strictly.
Quick Recall
What ΔEN range defines a polar covalent bond?
Why can period 3+ elements have an expanded octet but period 2 elements cannot?
Between ionic charge and ionic radius, which has the bigger effect on lattice energy?