Molecular shapes and bond angles are determined by the arrangement of electron pairs around a central atom, where repulsions between bonding and lone pairs control the three-dimensional geometry.

Molecular Geometry and the VSEPR Principle

The Valence Shell Electron Pair Repulsion (VSEPR) theory underpins the prediction of molecular shapes for A-Level Chemistry. It is based on the idea that electron pairs in the valence shell arrange themselves to minimise repulsion, producing specific molecular geometries and bond angles. This subsubtopic focuses on predicting shapes for molecules containing up to six electron pairs, recognising the influence of both bonding pairs and lone pairs.

VSEPR provides a systematic approach for identifying geometry by examining the number and type of electron pairs present. Lone pairs exert stronger repulsion than bonding pairs, causing characteristic deviations from idealised bond angles in many molecules.

You should be able to predict shapes and bond angles for molecules with up to six electron pairs around a central atom.

Diagram of the ideal electron‑pair geometries for 2–6 electron groups, illustrating linear, trigonal planar, tetrahedral, trigonal bipyramidal and octahedral shapes with their characteristic bond angles. Source

Electron Pair Repulsion and Its Consequences

Bonding Pairs vs Lone Pairs

Electron pair repulsion is not uniform. Different electron pair types influence molecular structure in specific ways:

• Lone pairs > bonding pairs in repulsive strength.

• Multiple bonds count as a single region of electron density despite containing more electrons.

• Increasing numbers of lone pairs generally reduce bond angles between bonding pairs.

Because lone pairs occupy more space, they compress bond angles and distort ideal shapes. This effect becomes especially important when comparing real molecular angles with theoretical predictions.

Predicting Shapes Using Electron Pair Numbers

Determining Regions of Electron Density

To apply VSEPR:

• Count the total number of regions of electron density (bonding pairs + lone pairs) around the central atom.

• Choose the geometry corresponding to this number.

• Adjust the name of the shape if lone pairs are present (e.g., trigonal pyramidal, bent, square pyramidal).

Between explanation and definition, the text should help students connect electron pair counts directly with real molecular geometry.

Ideal Geometries for 2–6 Electron Pairs

Two electron pairs

• Arrangement: Linear

• Typical bond angle: 180°

• Lone pairs: Uncommon in stable two-pair molecules.

Three electron pairs

• Arrangement: Trigonal planar

• Ideal angle: 120°

• With one lone pair: Bent shape, angle slightly <120°.

Four electron pairs

• Arrangement: Tetrahedral

• Ideal angle: 109.5°

• One lone pair: Trigonal pyramidal, angle ~107°.

• Two lone pairs: Bent, angle ~104.5°.

These variations demonstrate how lone pairs progressively compress bond angles.

Five electron pairs

• Arrangement: Trigonal bipyramidal

• Bond angles: 90°, 120°, 180°

• One lone pair: Seesaw shape.

• Two lone pairs: T-shaped.

• Three lone pairs: Linear.

Six electron pairs

• Arrangement: Octahedral

• Bond angles: 90°, 180°

• One lone pair: Square pyramidal.

• Two lone pairs: Square planar.

These models allow students to predict reliable shapes and angles using only electron pair counts, aligning directly with the OCR specification.

Lone Pair Effects on Bond Angles

Why Lone Pairs Distort Shapes

Lone pairs occupy electron density closer to the nucleus and are more compressed, increasing repulsion with bonding pairs. As a result:

• Bond angles decrease as lone pairs are added.

• The distortion follows a general pattern:
  Lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.

This hierarchy explains the well-known progression of decreasing angles across tetrahedral-based shapes.

Common Angle Deviations

Several bond angles frequently appear in OCR exams:

• 109.5° → 107° → 104.5° as lone pairs increase around a tetrahedral arrangement.

• 120° → slightly less in trigonal planar systems with one lone pair.

• 90° angles in octahedral systems remain generally robust but may subtly reduce with lone pair presence.

These values should be recognised and memorised, as they are fundamental for accurate molecular geometry predictions.

Shapes for Molecules with Six or Fewer Electron Pairs

Summary of Key Shapes (Bullet-Point Format)

• Linear (2 pairs)
  – 180° bond angle
  – Symmetrical electron distribution

• Trigonal planar (3 pairs)
  – 120°
  – Distorted to bent with one lone pair

• Tetrahedral (4 pairs)
  – 109.5°
  – Trigonal pyramidal (1 lone pair)
  – Bent (2 lone pairs)

• Trigonal bipyramidal (5 pairs)
  – 90°, 120°, 180°
  – Seesaw, T-shaped, linear (with increasing lone pairs)

• Octahedral (6 pairs)
  – 90°, 180°
  – Square pyramidal or square planar with lone pairs

These forms represent the full range of shapes required by the specification.

Applying VSEPR to Real Molecules

Understanding electron pair repulsion allows students to make precise qualitative predictions:

• Identify central atom and total electron pairs.

• Determine ideal geometry.

• Adjust expected bond angles if lone pairs are present.

• Name the resulting shape using standard terminology.

Between explanation and definition, the text should help students connect electron pair counts directly with real molecular geometry.

In the tetrahedral family, compare methane (CH₄), ammonia (NH₃) and water (H₂O).

Molecular structure of ammonia showing three N–H bonds in a trigonal pyramidal shape and a lone pair on nitrogen, illustrating lone‑pair repulsion reducing bond angles. Source

In a trigonal bipyramidal arrangement, three positions are equatorial (120° apart in a plane) and two are axial (above and below, 90° to the plane).

Ball‑and‑stick model of a trigonal bipyramidal molecule, showing three equatorial positions at 120° and two axial positions at 90°, supporting visualisation of bond‑angle differences. Source