Molecular Geometry, Polarity, and Hybridization (Ch. 11.1–11.8): VSEPR, Bond Angles, and Molecular Orbitals

Molecular Geometry & Polarity

Introduction

This section covers the fundamental principles of molecular geometry, electron domain theory (VSEPR), the effect of lone pairs, molecular polarity, and the basics of hybridization and molecular orbital theory. Understanding these concepts is essential for predicting the shapes, bond angles, and properties of molecules.

VSEPR Theory (Valence Shell Electron Pair Repulsion)

Basic Principles

• Electron groups (domains) around a central atom repel each other and arrange themselves to minimize repulsion (maximize separation).

• Each of the following counts as one electron group:

  • Single bond

  • Double bond

  • Triple bond

  • Lone pair

  • Single unpaired electron

• The preferred geometry is the one with the maximum separation possible (minimum energy).

Counting Electron Groups

• Example: The nitrite ion, $\mathrm{NO_2^-}$, has 3 electron groups on the central atom (2 bonding groups, 1 lone pair).

The Five Basic Electron Geometries

Number of Electron Domains	Electron-Domain Geometry	Predicted Bond Angles	Example
2	Linear	180°	CO2
3	Trigonal planar	120°	BF3
4	Tetrahedral	109.5°	CH4
5	Trigonal bipyramidal	120°, 90°	PF5
6	Octahedral	90°	SF6

Electron Geometry vs. Molecular Geometry

Electron geometry considers all electron groups (bonding and lone pairs), while molecular geometry considers only the arrangement of atoms (lone pairs are "invisible").

Electron Domains	Electron Geometry	0 Lone Pairs	1 Lone Pair	2 Lone Pairs	3 Lone Pairs	4 Lone Pairs
2	Linear	Linear
3	Trigonal planar	Trigonal planar	Bent
4	Tetrahedral	Tetrahedral	Trigonal pyramidal	Bent
5	Trigonal bipyramidal	Trigonal bipyramidal	Seesaw	T-shaped	Linear
6	Octahedral	Octahedral	Square pyramidal	Square planar

The Effect of Lone Pairs

• Linear: Molecular geometry is also linear.

• Trigonal planar: Can be trigonal planar (no lone pairs) or bent (1 lone pair).

• Tetrahedral: Can be tetrahedral (no lone pairs), trigonal pyramidal (1 lone pair), or bent (2 lone pairs).

• Trigonal bipyramidal: Can be trigonal bipyramidal (no lone pairs), seesaw (1 lone pair), T-shaped (2 lone pairs), or linear (3 lone pairs).

• Octahedral: Can be octahedral (no lone pairs), square pyramidal (1 lone pair), or square planar (2 lone pairs).

Electron Group Repulsion

• Lone pairs and single electrons occupy more space than bonding pairs, causing greater repulsion and smaller bond angles.

• Multiple bonds (double, triple) take up more space than single bonds, but not as much as lone pairs.

• Order of repulsion: Lone pair/single electron >> Triple bond > Double bond > Single bond

Examples of Electron and Molecular Geometry

• Linear: $\mathrm{BeCl_2}$ (linear, 180°)

• Trigonal planar: $\mathrm{BF_3}$ (trigonal planar, 120°), $\mathrm{SO_2}$ (bent, <120°)

• Tetrahedral: $\mathrm{CH_4}$ (tetrahedral, 109.5°), $\mathrm{NH_3}$ (trigonal pyramidal, <109.5°), $\mathrm{H_2O}$ (bent, <109.5°)

• Trigonal bipyramidal: $\mathrm{PCl_5}$ (trigonal bipyramidal, 120°, 90°), $\mathrm{SF_4}$ (seesaw, <120°, <90°), $\mathrm{ClF_3}$ (T-shaped), $\mathrm{XeF_2}$ (linear)

• Octahedral: $\mathrm{SF_6}$ (octahedral, 90°), $\mathrm{BrF_5}$ (square pyramidal), $\mathrm{XeF_4}$ (square planar)

Polarity of Molecules

Polar Bonds vs. Polar Molecules

• Polar covalent bonds have unequal sharing of electrons, resulting in a bond dipole.

• Molecular polarity depends on the vector sum of all bond dipoles (bond-dipole vector addition).

• Even if a molecule has polar bonds, the molecule may be nonpolar if the dipoles cancel (e.g., $\mathrm{CO_2}$).

• Examples:

  • $\mathrm{H_2O}$: Polar molecule (net dipole moment)

  • $\mathrm{CO_2}$: Nonpolar molecule (bond dipoles cancel)

Determining Molecular Polarity

• Linear: Nonpolar if terminal atoms are the same; polar if different.

• Trigonal planar: Nonpolar if terminal atoms are the same; bent is always polar due to lone pair.

• Tetrahedral: Nonpolar if terminal atoms are the same; trigonal pyramidal and bent are always polar due to lone pairs.

• Trigonal bipyramidal: Nonpolar if terminal atoms are the same; seesaw and T-shaped are always polar due to lone pairs.

• Octahedral: Nonpolar if terminal atoms are the same; square pyramidal is always polar due to lone pair; square planar is nonpolar if terminal atoms are the same.

Hybridization & Molecular Orbitals

Valence Bond Theory

• Atomic orbitals in a molecule combine to form hybrid orbitals that better represent the actual distribution of electrons in bonds.

• Hybrid orbitals have different shapes, orientations, and energies than standard atomic orbitals.

• Hybridization minimizes the energy of the molecule by maximizing orbital overlap.

Types of Hybridization

• sp3 hybridization: Tetrahedral electron geometry (e.g., CH4, NH3, H2O)

• sp2 hybridization: Trigonal planar electron geometry (e.g., BF3, SO2)

• sp hybridization: Linear electron geometry (e.g., BeCl2, CO2)

• sp3d and sp3d2 hybridization: Trigonal bipyramidal and octahedral geometries, respectively

Example: Methane ($\mathrm{CH_4}$)

• Has tetrahedral electron and molecular geometry

• Bond angle: 109.5°

• Uses four sp3 hybrid orbitals formed by mixing 2s, 2px, 2py, and 2pz atomic orbitals

Summary Table: Electron Geometry, Bond Angles, and Hybridization

Electron Geometry	Electron Groups	Bond Angles	Hybridization
Linear	2	180°	sp
Trigonal planar	3	120°	sp2
Tetrahedral	4	109.5°	sp3
Trigonal bipyramidal	5	120°, 90°	sp3d
Octahedral	6	90°	sp3d2

Key Takeaways

• Use VSEPR theory to predict molecular shapes and bond angles based on electron group repulsions.

• Distinguish between electron geometry (all electron groups) and molecular geometry (atoms only).

• Lone pairs and multiple bonds affect bond angles and molecular shape.

• Molecular polarity depends on both bond polarity and molecular geometry.

• Hybridization explains the observed shapes of molecules and the formation of equivalent bonds.