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S1.1 - Introduction to the particulate nature of matter S1.2 - The nuclear atom S1.3 - Electron configurations S1.4 - Counting particles by mass - The mole S1.5 - Ideal gases S2.1 - The ionic model S2.2 - The covalent model S2.3 - The metallic model S2.4 - From models to materials S3.1 - The periodic table - Classification of elements S3.2 - Functional groups - Classification of organic compounds R1.1 - Measuring enthalpy changes R1.2 - Energy cycles in reactions R1.3 - Energy from fuels R1.4 - Entropy and spontaneity AHL R2.1 - How much? The amount of chemical change R2.2 - How fast? The rate of chemical change R2.3 - How far? The extent of chemical change R3.1 - Proton transfer reactions R3.2 - Electron transfer reactions R3.3 - Electron sharing reactions R3.4 - Electron-pair sharing reactions

S2.2 - The covalent model

2.2.1 Covalent Bonds and Lewis Formulas 2.2.2 Bond Types 2.2.3 Co-coordination (Dative) Bonds 2.2.4 VSEPR Shapes of Molecules 2.2.5 Electronegativity and Bond Polarity 2.2.6 Polarity and Dipole Moments 2.2.7 Covalent Network Structures and Allotropes 2.2.8 Intermolecular Forces 2.2.9 Physical Properties of Covalent Substances 2.2.10 Chromatography and Intermolecular Forces 2.2.11 Resonance Structures (AHL) 2.2.12 Benzene and Resonance (AHL) 2.2.13 Expanded Octet and VSEPR (AHL) 2.2.14 Formal Charge (AHL) 2.2.15 Sigma and Pi Bonds (AHL) 2.2.16 Hybridization (AHL)

Shapes of molecules

Specification Reference S2.2.4

Quick Notes

  • Electron pairs around a central atom repel and arrange themselves as far apart as possible to minimise repulsion, giving different bonding shapes and arrangements.
  • Repulsion order: Lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.
  • The shape of a molecule is determined by the number of bonding and lone pairs around the central atom.
  • Common shapes and bond angles:
    • Linear (180°) – e.g. CO2
    • Bent (104°) – e.g. H2O
    • Trigonal planar (120°) – e.g. BF3
    • Trigonal pyramidal (107°) – e.g. NH3
    • Tetrahedral (109.5°) – e.g. CH4
    • Square planar (90°) – e.g. XeF4
    • Trigonal bipyramidal (90°, 120°, 180°) – e.g. PCl5
    • Octahedral (90°, 180°) – e.g. SF6

Full Notes

Electron Pair Repulsion Theory

The shape of a molecule can be predicted based on the number of bonds and lone pairs around the central atom. Bonds and lone pairs are considered regions of electron density.

Common Molecular Shapes and Bond Angles

Quick Reference Summary Table at Bottom of Page

Linear (180°)

2 bonding pairs, no lone pairs → bonds remain in a straight line.

CIE A-Level Chemistry diagram showing linear shape with 180° bond angle

Trigonal Planar (120°)

3 bonding pairs, no lone pairs → flat triangle arrangement.

Examples: BF3, NO3−

CIE A-Level Chemistry diagram showing trigonal planar shape with 120° bond angle

Tetrahedral (109.5°)

4 bonding pairs, no lone pairs → 3D tetrahedral shape.

Examples: CH4, NH4+

CIE A-Level Chemistry diagram showing tetrahedral shape with 109.5° bond angle

Trigonal Pyramidal (107°)

3 bonding pairs, 1 lone pair → bond angle reduced due to lone pair repulsion.

Examples: NH3, PCl3

AQA A-Level Chemistry diagram showing trigonal pyramidal shape with 107° bond angle

Bent (104.5°)

2 bonding pairs, 2 lone pairs → bond angle reduced further by two lone pairs.

Examples: H2O, OF2

CIE A-Level Chemistry diagram showing bent V-shaped structure with 104.5° bond angle

Trigonal Bipyramidal (90°, 120°, 180°)

5 bonding pairs, no lone pairs → atoms arranged in two layers.

Example: PCl5

CIE A-Level Chemistry diagram showing trigonal bipyramidal shape with bond angles of 90°, 120°, and 180°

Octahedral (90°, 180°)

6 bonding pairs, no lone pairs → symmetrical 3D shape.

Example: SF6

CIE A-Level Chemistry diagram showing octahedral shape with 90° and 180° bond angles

Square Planar (90°)

4 bonding pairs, 2 lone pairs → lone pairs opposite, minimising repulsion.

Example: XeF4

CIE A-Level Chemistry diagram showing square planar structure with 90° bond angles

Effect of Lone Pairs on Bond Angles

Lone pairs repel bonding pairs more than bonding pairs repel each other. This pushes bonding pairs closer together and reduces bond angles.

CIE A-Level Chemistry diagram showing how lone pairs affect bond angles
Lone Pairs Present Bond Angle Reduction Example
0 No reduction CH4 (109.5°)
1 ~2.5° smaller NH3 (107°)
2 ~5° smaller H2O (104.5°)

Molecular Shapes and Bond Angles – Key Examples

Molecule Electron Pair Geometry Shape Bond Angle(s) Explanation
CO2 2 bonding pairs Linear 180° No lone pairs, equal repulsion between bonds keeps atoms in a straight line
BF3 3 bonding pairs Trigonal planar 120° Bonds spread evenly in one plane with equal repulsion
CH4 4 bonding pairs Tetrahedral 109.5° Four bonds repel equally in 3D space, forming a symmetrical shape
NH3 3 bonding + 1 lone pair Pyramidal 107° Lone pair pushes bonding pairs slightly closer together
H2O 2 bonding + 2 lone pairs Non-linear 104.5° Two lone pairs create even more repulsion, reducing angle further
PF5 5 bonding pairs Trigonal bipyramidal 120° (eq), 90° (ax) Three bonds form a triangle in one plane; two others are perpendicular
SF6 6 bonding pairs Octahedral 90° All 6 electron pairs repel equally, forming a symmetrical 3D shape

Application in Ions

The same rules as above apply for polyatomic ions.

For Example:

Summary

Shape Bond Angle Lone Pairs? Example
Linear 180° No CO2
Trigonal Planar 120° No BF3
Tetrahedral 109.5° No CH4
Trigonal Pyramidal 107° 1 NH3
Bent (V‑Shaped) 104.5° 2 H2O
Trigonal Bipyramidal 90° & 120° No PCl5
Seesaw <90° & <120° 1 SF4
T‑Shaped <90° 2 ClF3
Octahedral 90° No SF6
Square Pyramidal <90° 1 BrF5
Square Planar 90° 2 XeF4