The Shapes of Simple Molecules & Ions (OCR A Level Chemistry A): Revision Note

Shapes of molecules

• The shape of a molecule depends on the number of bonding pairs and lone pairs of electrons around the central atom

• These electron pairs repel each other and determine the bond angles and geometry

Linear

• A linear shape occurs when there are two regions of electron density around the central atom

  • Typically from two double bonds or two single bonds with no lone pairs

• These bonding regions repel equally to minimise repulsion

• This creates a straight-line molecular geometry with a bond angle of 180^o

• Example: Carbon dioxide (CO₂), where the central carbon atom forms two double bonds and has no lone pairs

Trigonal planar

• A trigonal planar shape occurs when there are three bonding pairs and no lone pairs around the central atom

• The electron pairs repel equally and lie in the same plane

• This creates a flat, triangular shape with a bond angle of 120^o

• Example: Boron trifluoride (BF₃), where boron forms three covalent bonds and has an incomplete octet

Tetrahedral

• A tetrahedral shape forms when there are four bonding pairs and no lone pairs on the central atom

• The bonding pairs arrange themselves in 3D space to be as far apart as possible

• This gives equal bond angles of 109.5^o

• Example: Methane (CH₄), where carbon shares its four outer electrons with hydrogen atoms

Pyramidal

• A pyramidal shape forms when there are three bonding pairs and one lone pair on the central atom

• The lone pair repels the bonding pairs more strongly, pushing them slightly closer together

• This results in a bond angle of 107^o

• Example: Ammonia (NH₃), where nitrogen forms three bonds with hydrogen and retains one lone pair

Non-linear (bent)

• A non-linear or bent shape occurs when the central atom has two bonding pairs and two lone pairs

• The lone pairs exert strong repulsive forces, compressing the bond angle

• This results in a bond angle of 104.5^o

• Example: Water (H₂O), where oxygen forms two bonds and holds two lone pairs

Trigonal bipyramidal

• This shape arises when there are five bonding pairs and no lone pairs

• Three bonds lie in a horizontal plane at 120^o to each other, while two are positioned vertically at 90^o

• This arrangement minimises repulsion in three dimensions

• Example: Phosphorus pentachloride (PCl₅), where phosphorus bonds with five chlorine atoms

Octahedral

• An octahedral shape occurs when there are six bonding pairs and no lone pairs

• The bonding pairs repel equally in 3D space and arrange themselves at 90^o angles

• The result is a symmetrical octahedron

• Example: Sulfur hexafluoride (SF₆), where sulfur forms six bonds with fluorine atoms

Summary

• The most common molecular shapes and their associated bond angles are:

Electron pair repulsion

• The valence shell electron pair repulsion (VSEPR) theory explains how molecules adopt specific shapes and bond angles

• It states that electron pairs around a central atom:

  • Repel each other

  • Arrange themselves as far apart as possible to minimise repulsion

Key principles of VSEPR theory

• Valence electrons are the outer-shell electrons involved in bonding

• Electron pairs repel one another, with lone pairs exerting more repulsion than bonding pairs

• Molecules adopt the most stable 3D shape to minimise repulsion

• Double and triple bonds are treated as a single region of electron density, but repel slightly more than single bonds

Repulsion strengths

• Different types of electron pairs have different repulsive forces:

  • Lone pair – lone pair > lone pair – bond pair > bond pair – bond pair

• This is because lone pairs occupy more space and lie closer to the nucleus than bonding pairs