9.3 Chemical Bonding & VSEPR Molecular Geometry
Key Takeaways
- Ionic bonds transfer electrons (metal + nonmetal, large electronegativity difference ≳ 1.7); covalent bonds share electrons (two nonmetals); metallic bonds delocalize electrons across a cation lattice.
- Lewis structures place valence electrons to give every atom (except H) an octet; calculate formal charge = valence − (nonbonding + ½ bonding electrons) and minimize it across the structure.
- VSEPR geometry is set by electron domains: 2 → linear (180°), 3 → trigonal planar (120°), 4 → tetrahedral (109.5°), 5 → trigonal bipyramidal (90°/120°), 6 → octahedral (90°).
- Lone pairs compress bond angles below the electron-domain ideal (NH3 107°, H2O 104.5°); a molecule is polar if it has polar bonds AND an asymmetric shape.
- Electronegativity differences classify bonds: < 0.4 nonpolar covalent, 0.4–1.7 polar covalent, > 1.7 ionic (approximate thresholds).
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
| ΔEN | Bond character | Example |
|---|---|---|
| < 0.4 | Nonpolar covalent | H–H (0), C–H (0.4) |
| 0.4 – 1.7 | Polar covalent | O–H (1.4), C–O (1.0) |
| > 1.7 (approx.) | Ionic | Na–Cl (2.1), K–F (3.2) |
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:
- Count total valence electrons (sum across atoms, adjust for charge).
- Place the least electronegative atom at the center (H never central).
- Connect with single bonds (2 e− each).
- Complete octets on terminal atoms first.
- Move lone pairs to form multiple bonds if the central atom lacks an octet.
- 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
| Electron domains | Bonding / Lone pairs | Electron geometry | Molecular shape | Bond angle | Example |
|---|---|---|---|---|---|
| 2 | 2 / 0 | Linear | Linear | 180° | CO2, BeH2 |
| 3 | 3 / 0 | Trigonal planar | Trigonal planar | 120° | BF3, CO3²− |
| 3 | 2 / 1 | Trigonal planar | Bent | < 120° | SO2 |
| 4 | 4 / 0 | Tetrahedral | Tetrahedral | 109.5° | CH4, SO4²− |
| 4 | 3 / 1 | Tetrahedral | Trigonal pyramidal | ~107° | NH3 |
| 4 | 2 / 2 | Tetrahedral | Bent | ~104.5° | H2O |
| 5 | 5 / 0 | Trigonal bipyramidal | Trigonal bipyramidal | 90°/120° | PCl5 |
| 5 | 4 / 1 | Trigonal bipyramidal | See-saw | < 120°/< 90° | SF4 |
| 5 | 3 / 2 | Trigonal bipyramidal | T-shape | < 90° | ClF3 |
| 5 | 2 / 3 | Trigonal bipyramidal | Linear | 180° | XeF2 |
| 6 | 6 / 0 | Octahedral | Octahedral | 90° | SF6 |
| 6 | 5 / 1 | Octahedral | Square pyramidal | < 90° | BrF5 |
| 6 | 4 / 2 | Octahedral | Square planar | 90° | XeF4 |
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?
- Valence electrons: Xe (8) + 4 F (7 each) = 36 e−.
- Four Xe–F single bonds use 8 e−; 28 e− remain.
- Each F gets 6 e− (3 lone pairs) = 24 e−; 4 e− remain on Xe = 2 lone pairs.
- 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.
What is the molecular geometry and approximate bond angle of NH3?
Which of the following molecules is nonpolar despite containing polar bonds?