1.5.1 Shapes of Simple Molecules & Ions

Key Principles

Electron Pairs as Charge Clouds

• Bonding pairs and lone pairs (non-bonding) are charge clouds that repel each other.
• Electrons in the outer shell of atoms arrange themselves as far apart as possible to minimize repulsion.

Different Types of Repulsion

• Lone pair-lone pair repulsion is stronger than lone pair-bond pair repulsion.
• Lone pair-bond pair repulsion is stronger than bond pair-bond pair repulsion.
• These differences in repulsion affect bond angles and molecular shapes.

Effect on Bond Angles

• The bond angle decreases as the number of lone pairs increases due to their stronger repulsive forces.

The Valence Shell Electron Pair Repulsion (VSEPR) Theory

• This theory is used to predict the shape of molecules based on the idea that electron pairs around a central atom repel each other.
• Molecules take the shape that allows the electron pairs to be as far apart as possible, minimizing repulsion.

Common Molecular Shapes and Bond Angles

1. Linear Shape

• Example: $CO₂, BeCl₂$
• Electron Pairs: 2 bonding pairs, 0 lone pairs.
• Bond Angle: 180°.
• Explanation: The two bonding pairs repel each other, positioning themselves on opposite sides.

2. Trigonal Planar

• Example: $BF₃$
• Electron Pairs: 3 bonding pairs, 0 lone pairs.
• Bond Angle: 120°.
• Explanation: The bonding pairs are equally spaced around the central atom.

3. Tetrahedral

• Example: $CH₄, NH₄⁺$
• Electron Pairs: 4 bonding pairs, 0 lone pairs.
• Bond Angle: 109.5°.
• Explanation: The bonding pairs are arranged in a tetrahedral shape to minimize repulsion.

4. Trigonal Pyramidal

• Example: $NH₃$
• Electron Pairs: 3 bonding pairs, 1 lone pair.
• Bond Angle: 107°.
• Explanation: The lone pair repels more strongly, pushing the bonding pairs closer together.

5. Bent or V-Shaped

• Example: $H₂O$
• Electron Pairs: 2 bonding pairs, 2 lone pairs.
• Bond Angle: 104.5°.
• Explanation: Two lone pairs exert stronger repulsion, reducing the bond angle.

6. Trigonal Bipyramidal

• Example: $PCl₅$
• Electron Pairs: 5 bonding pairs, 0 lone pairs.
• Bond Angles: 120° (equatorial) and 90° (axial).
• Explanation: Bonding pairs arrange in two planes to minimize repulsion.

7. Octahedral

• Example: $SF₆$
• Electron Pairs: 6 bonding pairs, 0 lone pairs.
• Bond Angle: 90°.
• Explanation: The six bonding pairs arrange themselves symmetrically around the central atom.

8. Square Planar

• Example: $XeF₄$
• Electron Pairs: 4 bonding pairs, 2 lone pairs.
• Bond Angle: 90°.
• Explanation: The lone pairs are placed opposite each other, leading to a flat, square shape.

Lone Pair Influence on Molecular Shapes

• Lone pairs take up more space than bonding pairs because they are only attracted to one nucleus. This leads to greater repulsion and smaller bond angles.
• Example: In water (H₂O), the bond angle is 104.5° due to the two lone pairs, which repel the bonding pairs more strongly.

Summary of Shapes Based on Electron Pairs

Shape	Bonding Pairs	Lone Pairs	Bond Angle	Example
Linear	2	0	180°	$CO₂$
Trigonal Planar	3	0	120°	$BF₃$
Tetrahedral	4	0	109.5°	$CH₄$
Trigonal Pyramidal	3	1	107°	$NH₃$
Bent (V-Shaped)	2	2	104.5°	$H₂O$
Trigonal Bipyramidal	5	0	90° & 120°	$PCl₅$
Octahedral	6	0	90°	$SF₆$
Square Planar	4	2	90°	$XeF₄$