Study Notes

VSEPR (Valence Shell Electron Pair Repulsion) theory predicts the shape of molecules based on the idea that electron pairs repel each other and arrange themselves as far apart as possible.

Region of electron density: any group of electrons around a central atom that acts as a unit.

• A single bond = 1 region

• A double bond = 1 region (even though it has 4 electrons)

• A triple bond = 1 region

• A lone pair = 1 region

To determine geometry:

1. Draw the Lewis structure

2. Count total regions of electron density around the central atom

3. Determine electron geometry (based on total regions)

4. Determine molecular geometry (based on BONDED regions only — ignore lone pairs)

Key: Electron geometry = arrangement of ALL regions. Molecular geometry = arrangement of ATOMS only.

The electron geometry is determined by the TOTAL number of electron density regions. The molecular geometry is determined by the arrangement of ATOMS only.

2 regions → Electron geometry: LINEAR → Molecular geometry: LINEAR

Example: CO₂ (O=C=O, no lone pairs) → linear, 180°

3 regions → Electron geometry: TRIGONAL PLANAR

• 0 lone pairs → Molecular geometry: TRIGONAL PLANAR (120°)

• 1 lone pair → Molecular geometry: BENT (< 120°)

Examples: BF₃ (trigonal planar); SO₂ (bent)

4 regions → Electron geometry: TETRAHEDRAL

• 0 lone pairs → Molecular geometry: TETRAHEDRAL (109.5°)

• 1 lone pair → Molecular geometry: TRIGONAL PYRAMIDAL (< 109.5°)

• 2 lone pairs → Molecular geometry: BENT (< 109.5°)

Examples: CH₄ (tetrahedral), NH₃ (trigonal pyramidal), H₂O (bent)

Reference sheet notation: AX = bonded atom, E = lone pair on central atom.

"Linear AX₂, trigonal planar AX₃, bent AX₂E/AX₂E₂, tetrahedral AX₄, trigonal pyramidal AX₃E"

Hybridization explains how atomic orbitals mix to form the electron geometry of a molecule.

• 2 regions of electron density → sp hybridization (linear, 180°)

• 3 regions of electron density → sp² hybridization (trigonal planar, 120°)

• 4 regions of electron density → sp³ hybridization (tetrahedral, 109.5°)

Key insight: The hybridization is determined by the NUMBER OF REGIONS OF ELECTRON DENSITY (including lone pairs), not by the molecular geometry.

Examples:

• BeCl₂ (2 regions, linear): sp

• BF₃ (3 regions, trigonal planar): sp²

• CH₄ (4 regions, tetrahedral): sp³

• H₂O (4 regions, bent): sp³ (2 bonds + 2 lone pairs = 4 regions)

• NH₃ (4 regions, trigonal pyramidal): sp³

When atoms form bonds, the type of orbital overlap determines if it's a sigma or pi bond.

Sigma (σ) bond: direct, head-on overlap of orbitals between two atoms.

• Every single bond is 1 sigma bond

• The first bond in any double or triple bond is always sigma

Pi (π) bond: side-by-side overlap of p orbitals, above and below the bond axis.

• A double bond = 1 σ + 1 π

• A triple bond = 1 σ + 2 π

Counting bonds:

• Single bond: 1σ, 0π

• Double bond: 1σ + 1π

• Triple bond: 1σ + 2π

Example: CO₂ (O=C=O)

• 2 double bonds = 2σ + 2π → total: 2 sigma bonds, 2 pi bonds

Molecular polarity depends on two factors:

1. Whether individual bonds are polar (unequal sharing due to EN difference)

2. Whether the bond dipoles cancel each other due to geometry

Nonpolar molecule: symmetrical arrangement of identical bonds → dipoles cancel

Example: CO₂ (O=C=O, linear, bond dipoles point in opposite directions → cancel)

Polar molecule: unsymmetrical arrangement → dipoles don't cancel

Example: H₂O (bent, both O-H dipoles point toward O → don't cancel)

Intermolecular Forces (IMFs): attractions between molecules (not within).

• London Dispersion Forces (LDF): weakest, present in ALL molecules (polar and nonpolar). Caused by temporary dipoles. Larger/heavier molecules have stronger LDF.

• Dipole-dipole: between POLAR molecules. Stronger than LDF alone.

• Hydrogen bonding: strongest IMF. Occurs ONLY when H is bonded directly to N, O, or F (these are highly electronegative and small). The H on one molecule is attracted to the lone pair of N, O, or F on another molecule.

Comparing IMF strength:

• For similar polar molecules: more polarized bonds (larger EN difference) = stronger IMF

• For nonpolar molecules: larger molecule = more electrons = stronger LDF