9.3 Chemical Bonding & VSEPR Molecular Geometry

Three Bond Types

Quick Answer: Ionic bonds transfer electrons between a metal and a nonmetal; covalent bonds share electrons between two nonmetals; metallic bonds delocalize valence electrons across a lattice of metal cations. The electronegativity difference (ΔEN) between the bonded atoms is the practical classifier.

Electronegativity Classifiers

These are thresholds, not cliffs — bond character is continuous.

Ionic Bonding

An ionic compound is a lattice of alternating cations and anions held by Coulombic attraction. Lattice energy grows with ion charge magnitude and shrinks with ionic radius. NaCl: each Na+ is surrounded by 6 Cl− and vice versa. Ionic compounds have high melting points, conduct when molten or dissolved, and form crystalline solids.

Covalent Bonding

Atoms share electron pairs. A single bond is one shared pair (σ); a double bond is one σ + one π; a triple is one σ + two π. Bond order correlates with strength and inversely with length: C–C (154 pm) > C=C (134 pm) > C≡C (120 pm).

Metallic Bonding

Metal nuclei sit in a "sea" of delocalized valence electrons. This explains electrical/thermal conductivity, malleability, ductility, and the luster of metals. Alloys are mixtures within this model.

Lewis Structures

Steps:

1. Count total valence electrons (sum across atoms, adjust for charge).
2. Place the least electronegative atom at the center (H never central).
3. Connect with single bonds (2 e− each).
4. Complete octets on terminal atoms first.
5. Move lone pairs to form multiple bonds if the central atom lacks an octet.
6. Check formal charges; prefer the structure with charges closest to zero and any negative charge on the more electronegative atom.

Formal Charge

Formal charge = valence e− − (nonbonding e− + ½ bonding e−)

Worked example: the sulfate ion SO4²−. Total valence = 6 + 4(6) + 2 = 32 e−. With S central and four S–O single bonds, every O has an octet and S has an octet, but formal charges are S = +2, each O = −1 (sum = +2 − 4 = −2 ✓). Resonance with two S=O double bonds gives S = 0, two O = 0, two O = −1 (sum −2 ✓) — preferred because it lowers formal charge and obeys octet for second-period S (expanded octet allowed for third period and below).

VSEPR Molecular Geometry

Valence Shell Electron Pair Repulsion says electron domains around a central atom arrange to minimize repulsion. An electron domain is a single bond, a multiple bond (counts as ONE domain), or a lone pair.

VSEPR Shape Table

Lone-pair rule: lone pairs occupy more space than bonding pairs, compressing bond angles. That is why NH3 (107°) and H2O (104.5°) are below the tetrahedral 109.5°.

Polarity

A bond is polar if ΔEN ≳ 0.4. A molecule is polar if it has polar bonds AND the bond dipoles do not cancel:

• CO2 (linear, two equal opposite dipoles) → nonpolar.
• H2O (bent, two equal dipoles at 104.5°) → polar.
• CCl4 (tetrahedral, four equal dipoles) → nonpolar.
• CHCl3 (tetrahedral, three C–Cl dipoles dominate) → polar.

Worked VSEPR Example

What is the molecular geometry of XeF4?

1. Valence electrons: Xe (8) + 4 F (7 each) = 36 e−.
2. Four Xe–F single bonds use 8 e−; 28 e− remain.
3. Each F gets 6 e− (3 lone pairs) = 24 e−; 4 e− remain on Xe = 2 lone pairs.
4. 6 electron domains (4 bonding + 2 lone) → octahedral electron geometry. The two lone pairs go trans to minimize repulsion; the four F occupy a square plane → square planar, 90° F–Xe–F angles, nonpolar.

This is a classic PA-CAT-style geometry question: count domains, place lone pairs to minimize repulsion, name the resulting shape, and decide polarity.