Chemical Bonding: Ionic, Covalent, and Molecular Shapes

Why atoms bond: the octet rule

Most main-group atoms "want" 8 valence electrons (the octet) — the configuration of a noble gas. They achieve this by sharing, transferring, or pooling electrons with other atoms. The form this takes defines the bond type.

Ionic bonding — electron transfer

A metal LOSES electrons; a nonmetal GAINS them. Result: oppositely charged ions held together by electrostatic attraction. Classic example:

Na (loses 1 e⁻) → Na⁺      Cl (gains 1 e⁻) → Cl⁻      Na⁺ + Cl⁻ → NaCl

Ionic compound properties:

• High melting/boiling points (strong electrostatic forces)
• Hard but brittle (shifted ions repel)
• Conduct electricity when MOLTEN or DISSOLVED (free ions), not as solid
• Often soluble in polar solvents (water)

To balance charges in a formula: cross the charges. Ca²⁺ + Cl⁻ → CaCl₂. Al³⁺ + O²⁻ → Al₂O₃.

Covalent bonding — electron sharing

Two nonmetals share one or more pairs of electrons. Each shared pair = one bond. Three flavors:

• Single bond — 1 shared pair (2 electrons). Example: H-H in H₂.
• Double bond — 2 shared pairs (4 electrons). Example: O=O in O₂, or C=O in CO₂.
• Triple bond — 3 shared pairs (6 electrons). Example: N≡N in N₂. Very strong, short.

Order of strength and length: triple > double > single (strongest, shortest first).

Polar vs nonpolar covalent

If the two bonded atoms have equal electronegativity, electrons are shared EQUALLY (nonpolar covalent). If unequal, electrons are pulled toward the more electronegative atom → POLAR covalent, creating partial charges (δ⁺ and δ⁻).

Rules of thumb (using Pauling electronegativity differences):

• ΔEN < 0.5 → nonpolar covalent (e.g., C-H)
• 0.5 ≤ ΔEN ≤ 1.7 → polar covalent (e.g., H-O)
• ΔEN > 1.7 → ionic (e.g., Na-Cl)

Metallic bonding — pooled electrons

In a metal, valence electrons aren't held by individual atoms — they form a delocalized "sea" surrounding positive metal-ion cores. This explains metal properties:

• Conducts electricity — electrons flow freely
• Malleable, ductile — ions can shift past each other; the electron sea adjusts
• Lustrous — electrons absorb and re-emit visible light
• High thermal conductivity — electrons carry kinetic energy

Alloys are mixtures of metals (brass = Cu+Zn; bronze = Cu+Sn; steel = Fe+C) that retain metallic bonding but often improve mechanical properties.

Lewis structures — drawing the bonds

Lewis structures show valence electrons as dots and bonds as lines. Steps to build one:

1. Count total valence electrons from all atoms.
2. Place the LEAST electronegative atom at the center (usually).
3. Connect each peripheral atom to the center with a single bond.
4. Distribute remaining electrons as lone pairs to satisfy octets.
5. If atoms don't have full octets, convert lone pairs into double or triple bonds.

Example: CO₂. Total valence = 4 + 6 + 6 = 16. Two double bonds (O=C=O), each O with 2 lone pairs, C with no lone pairs.

VSEPR — predicting molecular shape

Valence Shell Electron Pair Repulsion theory: electron groups around a central atom arrange themselves as FAR APART AS POSSIBLE.

Lone pairs take up more space than bonded pairs, compressing the angle slightly.

Molecular polarity — geometry matters

A molecule is polar if it has polar bonds AND an asymmetric shape that prevents the bond dipoles from canceling. H₂O (bent) is polar — its bond dipoles add up. CO₂ (linear with identical bonded atoms) is NONpolar — the bond dipoles cancel exactly. CH₄ (tetrahedral, identical bonded atoms) is also nonpolar.

Intermolecular forces (IMF) — weaker but crucial

These are forces BETWEEN molecules (not within). From weakest to strongest:

• London dispersion — present in ALL molecules; temporary fluctuating dipoles.
• Dipole-dipole — between polar molecules.
• Hydrogen bonding — special case of dipole-dipole; H bonded to F, O, or N. Explains water's unusually high boiling point.

IMF strength determines boiling point, melting point, viscosity, surface tension — every "physical" property of liquids.

Check yourself

Practice with real CBE questions

• VSEPR molecular shape recognition (linear, bent, tetrahedral)
• Ionic vs covalent vs metallic bond classification