Shapes of Molecules

Learning Objectives (OCR Spec 2.2.2)

By the end of this lesson you should be able to:

• State and apply Valence Shell Electron Pair Repulsion (VSEPR) theory
• Predict the shape and bond angle of simple covalent molecules and ions
• Recognise the effect of lone pairs on molecular geometry
• Describe the shapes linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral, pyramidal and bent
• Explain why certain molecules have slightly reduced bond angles

VSEPR Theory: The Basic Idea

Valence Shell Electron Pair Repulsion (VSEPR) theory states that:

1. Electron pairs in the outer shell (valence shell) of a central atom repel one another
2. They arrange themselves as far apart as possible to minimise repulsion
3. The overall shape of the molecule is determined by the positions of the atoms (not lone pairs)
4. Lone pair repulsion is stronger than bond pair repulsion, so lone pairs "squeeze" bond pairs closer together

Repulsion hierarchy:

lone-lone > lone-bond > bond-bond

A double or triple bond is treated as a single "region of electron density" for VSEPR purposes (it points in one direction).

Electron Pair Geometries

Count the total number of electron pair regions (bond pairs + lone pairs) around the central atom. Then look at how many are lone pairs.

Regions	Electron geometry	Ideal angle
2	Linear	180 degrees
3	Trigonal planar	120 degrees
4	Tetrahedral	109.5 degrees
5	Trigonal bipyramidal	90 and 120 degrees
6	Octahedral	90 degrees

2 Regions: Linear (180 degrees)

• BeCl2 (only 2 bond pairs; incomplete octet): Cl-Be-Cl, 180 degrees
• CO2 (2 double bonds): O=C=O, 180 degrees
• HCN: H-C::N, 180 degrees

3 Regions: Trigonal Planar (120 degrees)

• BF3 (3 bond pairs, no lone pairs): T-shape of three F atoms in a plane, 120 degrees
• SO3: 3 S=O bonds arranged trigonally
• If 2 bond pairs + 1 lone pair -> bent shape ("V-shaped"), e.g. SO2 (bond angle ~117 degrees due to lone pair repulsion)

4 Regions: Tetrahedral (109.5 degrees)

This is the most common case in A-Level chemistry. Start with 4 bond pairs, then progressively replace bond pairs with lone pairs:

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

Each lone pair reduces the bond angle by roughly 2-2.5 degrees because it repels bond pairs more strongly than bond pairs repel each other.

5 Regions: Trigonal Bipyramidal

• PCl5 (5 bond pairs, 0 lone pairs): three equatorial Cl at 120 degrees, two axial Cl at 90 degrees to the plane
• SF4, ClF3, XeF2 have lone pairs in the equatorial positions, giving see-saw, T-shape and linear geometries respectively (beyond A-Level but useful context)

6 Regions: Octahedral (90 degrees)

• SF6 (6 bond pairs, 0 lone pairs): 6 F atoms around S, all at 90 degrees to each other
• XeF4 (4 bond pairs, 2 lone pairs): square planar
• [Fe(H2O)6]^3+ and other metal aquo complexes

Worked Examples

Example 1: NH3

• N has 5 outer electrons
• 3 shared with H atoms -> 3 bond pairs
• Remaining 2 on N -> 1 lone pair
• Total regions = 4 -> tetrahedral electron geometry
• With 1 lone pair -> trigonal pyramidal shape
• Bond angle ~107 degrees (reduced from 109.5 degrees by lone pair)

Example 2: H2O

• O has 6 outer electrons
• 2 shared with H atoms -> 2 bond pairs
• Remaining 4 on O -> 2 lone pairs
• Total regions = 4 -> tetrahedral electron geometry
• With 2 lone pairs -> bent / V-shaped
• Bond angle ~104.5 degrees (reduced from 109.5 degrees by 2 lone pairs)

Example 3: NH4+

• N has 4 bond pairs and NO lone pair (the original lone pair is now a dative bond to H+)
• 4 regions, 0 lone pairs -> perfect tetrahedral, bond angle 109.5 degrees

Example 4: CO2

• C has 2 C=O double bonds = 2 regions of electron density
• No lone pairs on C
• 2 regions -> linear, bond angle 180 degrees

Example 5: BF3

• B has only 3 bond pairs (incomplete octet)
• 3 regions, 0 lone pairs -> trigonal planar, bond angle 120 degrees

Example 6: SF6

• S has 6 bond pairs (expanded octet)
• 6 regions, 0 lone pairs -> octahedral, bond angle 90 degrees

Bond Angle Reductions due to Lone Pairs