Chemical Bonding II: Molecular Shapes, VSEPR Theory, and Electron Group Repulsions

Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory

Introduction to Molecular Shapes and Bonding Theories

Chemical bonding theories help us explain and predict the combinations of atoms that form stable molecules. The Lewis model provides a foundation for understanding molecular structure, while the Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the general shape of a molecule from its Lewis structure. Additional bonding theories, such as valence bond theory and molecular orbital theory, further expand our understanding of chemical bonding.

VSEPR Theory: The Five Basic Shapes

Fundamentals of VSEPR Theory

VSEPR theory is based on the principle that electron groups—including lone pairs, single bonds, multiple bonds, and even single electrons—repel one another through coulombic forces. This repulsion determines the geometry of the molecule.

• Electron group: Any region of electron density around a central atom, including lone pairs, single, double, or triple bonds, and single electrons.

• Each electron group counts as one domain for geometry determination, regardless of whether it is a lone pair or any type of bond.

• Repulsions between electron groups dictate the arrangement of atoms in space.

Example: In CO2, the two double bonds each count as one electron group, resulting in a linear geometry.

How to Determine Molecular Shapes

1. Draw the Lewis Structure of the molecule.

2. Count the number of electron groups around the central atom (lone pairs, single, double, triple bonds, or single electrons).

3. Consult the VSEPR table to determine the electron geometry and molecular shape.

Example: For CH4, there are four electron groups (four single bonds), resulting in a tetrahedral geometry.

Electron Group Geometry vs. Molecular Geometry

Distinguishing Electron Geometry and Molecular Geometry

The electron geometry describes the spatial arrangement of all electron groups (bonding and lone pairs) around the central atom. The molecular geometry describes the arrangement of only the atoms (excluding lone pairs).

• Electron geometry determines the overall scaffold of the molecule.

• Molecular geometry is the actual shape formed by the atoms.

Example: In NH3, the electron geometry is tetrahedral (four groups), but the molecular geometry is trigonal pyramidal (three atoms and one lone pair).

Effect of Lone Pairs on Molecular Geometry

Lone Pair Repulsion and Bond Angles

Lone pairs occupy more space than bonding pairs, leading to greater repulsion and smaller bond angles than the ideal geometry. This effect is especially pronounced in molecules with lone pairs on the central atom.

• Order of repulsion: Lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair

• Bond angles decrease as the number of lone pairs increases.

Example: In H2O, the electron geometry is tetrahedral, but the molecular geometry is bent due to two lone pairs, resulting in a bond angle of approximately 104.5°.

Common Electron Group Geometries and Molecular Shapes

Summary Table of Electron Group and Molecular Geometries

The following table summarizes the relationship between the number of electron groups, bonding groups, lone pairs, electron geometry, molecular geometry, approximate bond angles, and examples:

Electron Groups	Bonding Groups	Lone Pairs	Electron Geometry	Molecular Geometry	Approximate Bond Angles	Examples
2	2	0	Linear	Linear	180°	CO2
3	3	0	Trigonal planar	Trigonal planar	120°	BF3
3	2	1	Trigonal planar	Bent	<120°	SO2
4	4	0	Tetrahedral	Tetrahedral	109.5°	CH4
4	3	1	Tetrahedral	Trigonal pyramidal	~107°	NH3
4	2	2	Tetrahedral	Bent	~104.5°	H2O
5	5	0	Trigonal bipyramidal	Trigonal bipyramidal	120°, 90°	PCl5
5	4	1	Trigonal bipyramidal	Seesaw	<120°, <90°	SF4
5	3	2	Trigonal bipyramidal	T-shaped	<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

Key Points and Equations

• Electron group count: Determined from the Lewis structure; resonance structures may be used to determine the number of electron groups.

• Each of the following counts as a single electron group: Lone pair, single bond, double bond, triple bond, or single electron.

• Geometry is determined by electron group repulsions: Lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.

• Bond angles: Can deviate from ideal values due to the presence of lone pairs and multiple bonds.

Equation for number of electron groups:

$\text{Number of electron groups} = \text{Number of bonds} + \text{Number of lone pairs}$

Ideal bond angles:

• Linear: $180^\circ$

• Trigonal planar: $120^\circ$

• Tetrahedral: $109.5^\circ$

• Trigonal bipyramidal: $120^\circ$ (equatorial), $90^\circ$ (axial)

• Octahedral: $90^\circ$

Practice and Application

• Draw Lewis structures to identify electron groups.

• Apply VSEPR theory to predict molecular shapes.

• Consider the effect of lone pairs and multiple bonds on bond angles and geometry.

Example: For NO3-, use resonance structures to determine the number of electron groups around the central atom.

Summary of VSEPR Theory

• The shape of a molecule is determined by the repulsions among all electron groups on the central atom.

• Lone pairs exert greater repulsion than bonding pairs, leading to deviations from ideal bond angles.

• VSEPR theory provides a systematic approach to predicting molecular geometry from Lewis structures.

Additional info: The notes focus on VSEPR theory and do not cover valence bond theory or molecular orbital theory in detail. For a complete understanding, students should consult additional resources on these topics.