Molecular and Ionic Compound Structure — AP Chemistry
1. Lewis Structures and Formal Charge ★★☆☆☆ ⏱ 5 min
The octet rule states that most main-group atoms tend to form bonds to reach 8 valence electrons, but there are three key exceptions: hydrogen only needs 2 valence electrons, period 3 and higher elements can have expanded octets (>8 electrons) due to available d-orbitals, and radical species have an odd number of total valence electrons.
FC = V - \left(LP + \frac{1}{2}BP\right)
Where $V$ = number of valence electrons of the neutral free atom, $LP$ = number of lone pair electrons on the atom in the structure, and $BP$ = number of bonding electrons attached to the atom. The most stable resonance structure has formal charges closest to 0, with any negative formal charges located on the most electronegative atoms.
2. VSEPR Theory and Molecular Geometry ★★★☆☆ ⏱ 4 min
Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the 3D shape of molecules based on the principle that negatively charged electron domains (bonding pairs, where single/double/triple bonds all count as 1 domain, plus lone pairs) repel each other to maximize the distance between them.
Lone pair repulsion is stronger than bonding pair repulsion, following the order: $LP-LP > LP-BP > BP-BP$. This causes actual bond angles to be smaller than the ideal EDG angle when lone pairs are present.
3. Orbital Hybridisation ★★★☆☆ ⏱ 4 min
Hybridisation describes the mixing of pure atomic orbitals (s, p, d) to form identical hybrid orbitals that match the electron domain geometry of the central atom. This explains why bonds to identical terminal atoms have equal energy and bond length. Hybridisation correlates directly to the number of electron domains around the central atom:
Hybrid orbitals only form sigma ($\sigma$) bonds: this includes all single bonds, and the first bond in any double or triple bond. Remaining unhybridized p or d orbitals form pi ($\pi$) bonds, which are the second bond in a double bond and the third bond in a triple bond.
4. Molecular Polarity ★★★☆☆ ⏱ 3 min
A covalent bond is polar if the electronegativity difference between the two atoms is greater than 0.4, creating a dipole moment (a partial positive charge on the less electronegative atom, and a partial negative charge on the more electronegative atom). A molecule is polar overall only if two conditions are met:
- It has at least one polar covalent bond
- It has asymmetric molecular geometry that prevents individual bond dipoles from canceling out
Symmetric geometries (e.g. linear with two identical terminal atoms, tetrahedral with four identical terminal atoms) will always cancel individual bond dipoles, resulting in a nonpolar molecule even with polar bonds.
5. Intermolecular Forces ★★★☆☆ ⏱ 4 min
Intermolecular forces (IMFs) are attractive forces between separate molecules, much weaker than intramolecular covalent or ionic bonds. They determine key physical properties including boiling point, melting point, solubility, and vapor pressure. IMFs are ranked from strongest to weakest below:
- **Ion-dipole**: Attraction between a charged ion and a polar molecule, the strongest IMF, found in solutions of ionic compounds (e.g. $Na^+$ and water molecules in salt water)
- **Hydrogen bonding**: A special strong dipole-dipole attraction that only occurs when a hydrogen atom is covalently bonded directly to N, O, or F, and interacts with a lone pair on a separate N/O/F atom. Responsible for water's unusually high boiling point.
- **Dipole-dipole**: Attraction between the partial positive end of one polar molecule and the partial negative end of another polar molecule.
- **London Dispersion Forces (LDF)**: Temporary attraction from random electron cloud fluctuations that create temporary dipoles. Present in all molecules, only IMF in nonpolar molecules and noble gases. Strength increases with molar mass/number of electrons, as larger electron clouds are more polarizable.
A key rule for solubility is *like dissolves like*: polar solutes dissolve in polar solvents, and nonpolar solutes dissolve in nonpolar solvents. This is because solute-solvent intermolecular attractions must be strong enough to overcome original solute-solute and solvent-solvent attractions to form a solution.
Common Pitfalls
Why: Students memorize the octet rule as an absolute rule with no exceptions
Why: Students skip counting lone pairs and only look at bonding atoms to determine shape
Why: Students ignore geometric symmetry that cancels out individual bond dipole moments
Why: Students confuse number of bonds with number of electron domains
Why: Students see hydrogen in a molecule and automatically assume hydrogen bonding exists