Shapes of Molecules and VSEPR Theory

This lesson covers the shapes of molecules and ions using Valence Shell Electron Pair Repulsion (VSEPR) theory. You will learn to predict and explain molecular geometry, including the effects of lone pairs on bond angles. This is a key topic in AQA specification 3.1.3.

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VSEPR Theory: The Principle

Key Definition: VSEPR (Valence Shell Electron Pair Repulsion) theory states that electron pairs around a central atom arrange themselves as far apart as possible to minimise repulsion.

The shape of a molecule is determined by the number of electron pairs (both bonding and lone pairs) around the central atom. Electron pairs repel each other because they are all negatively charged.

Order of Repulsion Strength

Lone pair–lone pair  >  Lone pair–bonding pair  >  Bonding pair–bonding pair

Lone pairs are held closer to the central atom (attracted to one nucleus only), so they occupy more space and exert greater repulsion. This causes bond angles to be reduced from the ideal values.

Exam Tip: When explaining bond angles, always state the order of repulsion: LP–LP > LP–BP > BP–BP. Then explain how this affects the bond angle compared to the ideal geometry.

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Shapes with No Lone Pairs on the Central Atom

2 Bonding Pairs → Linear (180°)

Example: BeCl₂

    Cl — Be — Cl
         180°

• Be has 2 bonding pairs, 0 lone pairs
• The 2 electron pairs repel equally and arrange themselves at 180°
• Bond angle: 180°

Other examples: CO₂ (O═C═O, 180°), HCN (H—C≡N, 180°)

3 Bonding Pairs → Trigonal Planar (120°)

Example: BF₃

          F
          |
    F — B — F
       120°

• B has 3 bonding pairs, 0 lone pairs
• The 3 electron pairs arrange at 120° in a flat plane
• Bond angle: 120°

Other examples: AlCl₃ (monomer), CO₃²⁻, NO₃⁻

4 Bonding Pairs → Tetrahedral (109.5°)

Example: CH₄

           H
           |
      H — C — H
          |
          H
    (three-dimensional: 109.5°)

• C has 4 bonding pairs, 0 lone pairs
• The 4 electron pairs arrange tetrahedrally
• Bond angle: 109.5°

Other examples: CCl₄, SiH₄, NH₄⁺, BF₄⁻

5 Bonding Pairs → Trigonal Bipyramidal (90° and 120°)

Example: PCl₅

        Cl (axial)
        |
   Cl — P — Cl  (equatorial, 120° apart)
        |
        Cl (axial)
   + one more Cl equatorial

• P has 5 bonding pairs, 0 lone pairs
• 3 equatorial positions at 120° to each other
• 2 axial positions at 90° to equatorial
• Bond angles: 120° (equatorial) and 90° (axial-equatorial)

6 Bonding Pairs → Octahedral (90°)

Example: SF₆

        F
        |
   F — S — F
       /|\
      F  F  F

• S has 6 bonding pairs, 0 lone pairs
• All bond angles: 90°

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Shapes with Lone Pairs on the Central Atom

3 Electron Pairs (1 Lone Pair) → Bent/V-shaped

Example: SO₂

       ••
    O═S═O
     ~119°

• S has 2 bonding regions (double bonds count as one region) and 1 lone pair
• Based on trigonal planar arrangement but one position is a lone pair
• Bond angle: approximately 119° (less than 120° due to lone pair repulsion)

4 Electron Pairs (1 Lone Pair) → Pyramidal

Example: NH₃

       ••
       N
      /|\
     H H H
      107°

• N has 3 bonding pairs and 1 lone pair (4 pairs total)
• Based on tetrahedral arrangement but one position is a lone pair
• The lone pair compresses the H—N—H bond angles
• Bond angle: 107° (reduced from 109.5° by lone pair repulsion)

Other examples: PCl₃ (~107°), H₃O⁺ (~107°), NF₃ (~102°)

4 Electron Pairs (2 Lone Pairs) → Bent/V-shaped

Example: H₂O

      •• ••
        O
       / \
      H   H
      104.5°

• O has 2 bonding pairs and 2 lone pairs (4 pairs total)
• Based on tetrahedral arrangement but two positions are lone pairs
• Two lone pairs compress the bond angle more than one
• Bond angle: 104.5° (reduced from 109.5°)