Revise: Chemical Bonding and Intermolecular Forces
How and why atoms bond — ionic, covalent, and metallic bonding, electronegativity and bond polarity, and the intermolecular forces (van der Waals, hydrogen bonding) that act between molecules.
Atoms bond to reach 8 valence electrons (an octet), matching a noble gas.
Na loses 1 electron; Cl gains 1 — both reach a stable octet.
ΔEN < 0.4 nonpolar covalent, 0.4–1.7 polar covalent, > 1.7 ionic.
Na (0.93) and Cl (3.16): ΔEN = 2.23 → ionic.
Ionic bonds: complete electron transfer, metal → nonmetal.
Mg loses 2e⁻, O gains 2e⁻ → MgO.
Covalent bonds: atoms share electron pairs instead of transferring them.
Two O atoms share 2 pairs → O₂ (a double bond).
H bonded to N, O, or F creates an unusually strong intermolecular attraction.
Hydrogen bonds between water molecules give water its high boiling point.
Only outer-shell (valence) electrons take part in bonding.
Chlorine has 17 electrons total, but only 7 are valence electrons.
Atoms gain, lose, or share electrons to reach a stable 8-electron outer shell.
Mg forms Mg²⁺ by losing both valence electrons.
ΔEN = |EN₁ − EN₂| predicts ionic vs. polar covalent vs. nonpolar covalent.
C–H: ΔEN = 0.35 → essentially nonpolar.
Ionic compounds: high melting point, hard but brittle, conduct only molten/dissolved.
Solid NaCl does not conduct; dissolved NaCl does.
More shared electron pairs → shorter, stronger bonds.
N₂'s triple bond makes nitrogen gas very unreactive.
The more electronegative atom in a bond becomes δ−; the other becomes δ+.
In C=O, oxygen is δ− and carbon is δ+.
Temporary, fluctuating dipoles create weak attractions between ALL molecules.
Van der Waals forces alone let nonpolar O₂ condense into a liquid when cold enough.
Metal ions in a lattice + a delocalized "sea" of electrons.
Copper's delocalized electrons make it an excellent electrical conductor.