Shapes of Molecules

Spec Mapping — OCR H432 Module 2.2.2 — Bonding and structure, covering Valence Shell Electron Pair Repulsion (VSEPR) theory, prediction of molecular shapes and bond angles from the count of bond pairs and lone pairs around a central atom, the effect of lone pair repulsion on bond angles, and the canonical molecular geometries — linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral, trigonal pyramidal, bent, and the lone-pair-distorted variants (refer to the official OCR H432 specification document for exact wording).

VSEPR theory — Valence Shell Electron Pair Repulsion — is the workhorse predictive tool of A-Level molecular geometry. It treats the molecule as a central atom surrounded by regions of electron density (bonding pairs and lone pairs) that repel each other, and predicts that those regions arrange in 3D space to be as far apart as possible. From just the number of bond pairs and lone pairs around a central atom you can predict, to within a few degrees, the bond angles of nearly every simple molecule in the OCR specification. The theory was systematised by Sidgwick and Powell (1940) and refined by Gillespie and Nyholm (1957) into the form you study today. This lesson teaches the protocol, the seven canonical shapes (linear, trigonal planar, bent (3-region), tetrahedral, trigonal pyramidal, bent (4-region), trigonal bipyramidal, octahedral, square planar), the lone-pair compression effect on bond angles, and the worked examples that OCR examines most often.

Key Definitions:

• VSEPR — Valence Shell Electron Pair Repulsion theory: outer-shell electron pairs repel; arrange to maximise separation.
• Region of electron density — a single bond, double bond, triple bond, or lone pair (each counted as one "group" for VSEPR).
• Bond angle — the angle measured at the central atom between two adjacent bonds.
• Electron geometry — the arrangement of all electron pairs (bond + lone).
• Molecular geometry — the arrangement of the atoms only (lone pairs not visible).

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The four VSEPR rules

1. Electron pairs in the outer (valence) shell of a central atom repel one another (Coulombic repulsion + Pauli exchange).
2. They arrange to be as far apart as possible — minimising repulsion gives the most stable geometry.
3. The shape of the molecule is determined by the positions of the atoms — lone pairs are present but not visible in molecular geometry.
4. Lone-pair repulsion is stronger than bond-pair repulsion — lone pairs occupy more space because they are held by only one atom rather than shared between two.

Repulsion hierarchy:

$\text{lone pair} \leftrightarrow \text{lone pair} > \text{lone pair} \leftrightarrow \text{bond pair} > \text{bond pair} \leftrightarrow \text{bond pair}$lone pair↔lone pair>lone pair↔bond pair>bond pair↔bond pair

Crucial counting rule: a multiple bond (double or triple) is treated as a single region of electron density for VSEPR purposes — all the bonding electrons in a multiple bond point in the same direction from the central atom.

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Counting regions of electron density

flowchart TD
    A[Draw dot-and-cross for the species] --> B[Count bond pairs at central atom]
    B --> C[Count lone pairs at central atom]
    C --> D[Bond pairs + lone pairs = total regions]
    D --> E{Total regions?}
    E -- "2" --> F[Electron geometry: linear]
    E -- "3" --> G[Electron geometry: trigonal planar]
    E -- "4" --> H[Electron geometry: tetrahedral]
    E -- "5" --> I[Electron geometry: trigonal bipyramidal]
    E -- "6" --> J[Electron geometry: octahedral]
    F --> K[Subtract lone pairs to get molecular shape]
    G --> K
    H --> K
    I --> K
    J --> K
    K --> L[Apply lone-pair compression to predict bond angle]

Total regions	Electron geometry	Ideal bond angle
2	Linear	180°
3	Trigonal planar	120°
4	Tetrahedral	109.5°
5	Trigonal bipyramidal	90° and 120°
6	Octahedral	90°

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The seven canonical molecular shapes

Linear (2 BP, 0 LP) 180° e.g. BeCl₂, CO₂

Trigonal planar (3 BP, 0 LP) 120° e.g. BF₃, SO₃

Tetrahedral (4 BP, 0 LP) 109.5° e.g. CH₄, NH₄⁺

Trigonal pyramidal (3 BP, 1 LP) ·· ~107° e.g. NH₃, H₃O⁺, PCl₃

Bent (2 BP, 2 LP) ·· ·· ~104.5° e.g. H₂O, OF₂

Bent — 3 region (2 BP, 1 LP) ·· ~117° e.g. SO₂, O₃

Trigonal bipyramidal (5 BP) 90°, 120° e.g. PCl₅

Octahedral (6 BP, 0 LP) 90° e.g. SF₆, [Fe(H₂O)₆]³⁺

Square planar (4 BP, 2 LP) ·· ·· 90° e.g. XeF₄, [Ni(CN)₄]²⁻

Purple = central atom; green = bonded atom; ·· = lone pair on central atom; dashed = behind plane

2 regions — linear (180°)

• BeCl₂ (g) — 2 bond pairs, 0 lone pairs; incomplete octet. Cl–Be–Cl angle 180°.
• CO₂ — two C=O double bonds count as 2 regions; O=C=O angle 180°.
• HCN — H–C≡N; one single bond, one triple bond; 180°.

3 regions — trigonal planar (120°)

• BF₃ — 3 bond pairs, 0 lone pairs; incomplete octet on B. F–B–F angle 120°.
• SO₃ — three S=O double bonds (each counts as 1 region); trigonal planar; 120°.
• 3 region, 1 lone pair — bent: SO₂ has 2 S=O bonds + 1 lone pair on S; O–S–O angle ≈ 117° (less than ideal 120° because lone pair compresses).

4 regions — tetrahedral (109.5°)

The most common case at A-Level. The geometry depends on how many of the 4 regions are lone pairs:

Bond pairs	Lone pairs	Shape	Bond angle	Example
4	0	Tetrahedral	109.5°	CH₄, NH₄⁺, SiF₄, SO₄²⁻
3	1	Trigonal pyramidal	~107°	NH₃, H₃O⁺, PCl₃
2	2	Bent (V-shaped)	~104.5°	H₂O, OF₂

Each lone pair compresses the bond angle by roughly 2-2.5° because lone pair repulsion exceeds bond pair repulsion.

5 regions — trigonal bipyramidal (90° and 120°)

• PCl₅ (g) — 5 bond pairs, 0 lone pairs (expanded octet on P). Three equatorial Cl at 120°; two axial Cl at 90° to the equatorial plane. The axial bond is longer than the equatorial because of stronger 90° repulsions from the equatorial bonds.
• Lone pairs in trigonal bipyramidal always go to equatorial positions (more room than axial):
  • SF₄ — 4 BP + 1 LP → see-saw.
  • ClF₃ — 3 BP + 2 LP → T-shape.
  • I₃⁻ — 2 BP + 3 LP → linear.
• (See-saw, T-shape, and the I₃⁻ case are beyond OCR A-Level scope but are useful context.)

6 regions — octahedral (90°)

• SF₆ — 6 bond pairs, 0 lone pairs (expanded octet). All F–S–F angles 90° (180° across the molecule).
• XeF₄ — 4 bond pairs + 2 lone pairs (octahedral electron geometry, square-planar molecular shape; the two lone pairs sit trans-axially to minimise lone-lone repulsion).
• [Fe(H₂O)₆]³⁺ — 6 dative bonds from H₂O; octahedral metal complex (Year 13 transition-metal chemistry).

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The lone pair compression effect — the canonical tetrahedral series

The series CH₄, NH₃, H₂O is the cleanest demonstration of lone-pair compression:

Molecule	Bond pairs	Lone pairs	Shape	H–X–H bond angle
CH₄	4	0	Tetrahedral	109.5°
NH₃	3	1	Trigonal pyramidal	107°
H₂O	2	2	Bent	104.5°

Each lone pair occupies a slightly larger angular region than a bond pair because:

1. A lone pair is held by only one nucleus (the central atom), whereas a bond pair is shared between two nuclei.
2. The lone pair therefore extends further from the central atom in angular space, "fanning out" and exerting stronger angular repulsion on neighbouring electron pairs.

The net effect is that bond pairs are squeezed closer together, reducing the bond angle below the ideal tetrahedral 109.5°. The reduction is roughly 2–2.5° per lone pair — a useful rule of thumb for predicting bond angles when no exam value is given.

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Worked examples

Example 1 — NH₃