The shapes of molecules and their bond angles (the angles between the covalent bonds) are not random, but are determined according to a set of principles called Valence Shell Electron Pair Repulsion Theory. (VSEPR)
We will consider a simple molecule as consisting of a central atom, around which a number of other atoms are arranged. The outer shell of this central atom will contain electrons which are part of bonding pairs, and may also contain lone pairs. It is the mutual repulsion between these electron pairs which gives the molecule its shape and bond angles.
VSEPR Theory states: i) electron pairs repel to get as far apart as possible and ii) lone pairs repel more strongly than bonding pairs.
As a consequence of this, a lone pair will push bonding pairs closer together, decreasing the angle between them by 2.5º per lone pair.
Each specific shape has a name that reflects its geometry:
To understand how these shapes arise, we need to look at specific molecules and in particular to examine the number and type of electron pairs around the central atom.
Molecules with only bonding pairs around the central atom
e.g. BeCl2 Be has 2 bonding pairs Be has 0 lone pairs
The two bonding pairs repel as far as possible from one another, giving a bond angle of 180°. The lone pairs on Cl have no effect on the shape. The molecule is LINEAR.
e.g. BF3 B has 3 bonding pairs B has 0 lone pairs
There are 3 bonding pairs arranging themselves as far away as possible. This leads to 120° bond angles in a planar molecule. The shape is TRIGONAL PLANAR.
e.g. CH4 C has 4 bonding pairs C has 0 lone pairs
There are 4 bonding pairs arranging themselves as far away as possible. This leads to a 109.5° bond angle and a TETRAHEDAL shape.
e.g. PF5 P has 5 bonding pairs P has 0 lone pairs
There are 5 bonding pairs arranging themselves as far apart as possible. Three lie in a plane, with 120° bond angles and the other two are at 90° to this plane. The shape is TRIGONAL BIPYRAMIDAL.
e.g. SF6 S has 6 bonding pairs S has 0 lone pairs
There are 6 bonding pairs arranging themselves as far apart as possible. Each lies at 90° to the others. The shape is OCTAHEDRAL.
Molecules where the central atom has lone pairs as well as bonding pairs
The electron pairs will still arrange themselves to get as far apart as possible: both the lone pairs and the bonding pairs will do this. We then need to name the shape. The lone pairs behave as invisible forces pushing the bonding pairs into new positions, and in naming the shape we only consider where the bonding pairs have ended up under the influence of these invisible lone pairs. When we determine the bond angles we also need to remember that the repulsion of the lone pairs towards bonding pairs is stronger than the repulsion between bonding pairs, so bond angles can be reduced by 2.5º.
e.g. NF3 N has 3 bonding pairs N has 1 lone pair
The four electron pairs take up a tetrahedral arrangement to get as far apart as possible, so the bond angles are 109.5 – 2.5 = 107º. The shape the three bonding pairs take up around the central N atom is PYRAMIDAL.
e.g. H2S S has 2 bonding pairs S has 2 lone pairs
The four electron pairs take up a tetrahedral arrangement to get as far apart as possible, to the bond angles are 109.5 – (2 x 2.5) = 104.5º. The shape the two bonding pairs take up around the central S atom is NON-LINEAR or BENT.
Molecules with double covalent bonds
For the purposes of VSEPR Theory, a double bond acts exactly the same as a single bond, so we count it as one bonding pair, rather than two.
e.g. CO2 C has two bonding pairs C has 0 lone pairs
We count the central C atom as having two (rather than four) bonding pairs, counting each double bond as a single bonding pair, so the bond angle is 180º and the shape is LINEAR.
Shapes of Compound Ions
Compound ions (such as ammonium NH4+, nitrate NO3–, sulphate SO42- etc.) have an overall charge and participate in ionic bonding with oppositely charged ions, but they have covalent bonding inside the ion, and follow the same rules with regard to VESPR Theory to determine the shape of the ion.
e.g. ammonium NH4+ N has 4 bonding pairs and 0 lone pairs
The four bonding pairs arrange themselves to get as far apart as possible. There is no difference between a dative bond and a covalent bond once it has been formed, so the bond angles are 109.5º and the shape is TETRAHEDRAL.
We can summarise the relationship between the number and types of electron pairs around the central atom and the resulting shape and bond angles as shown in the table below. Once the number of lone pairs and bonding pairs are known, the appropriate shape and bond angles can be selected. (Here the lone pairs have been shown in grey so that their influence on the shape can be visualised, but remember the lone pairs aren’t part of the named shape).
Explaining the shapes of molecules
It is common in exam questions to be asked to state and explain the shape of a molecule. Having determined the name of the shape and the corresponding bond angles, we need to be able to explain our logic.
We need to include the following points:
- A statement of the number of bonding and lone pairs (even when none) around the central atom in the molecule we are discussing
- That electron pairs repel to get as far apart as possible
- That lone pairs repel more strongly than bonding pairs (if lone pairs are present)
- A statement of the shape the molecule has as a result
Representing the shapes of molecules
Drawing a 3D shape in 2D requires a convention to show when bonds are not in the plane of the paper. When drawing molecules, a bond drawn as a solid wedge means a bond coming forward out of the plane of the paper, and a bond draw as a dashed wedge means a bond going backwards behind the paper.
A methane molecule with tetrahedral shape, drawn to show its 3D arrangement, with the two bonds on the right of the molecule coming out of and into the paper, and the bond to the left and above the carbon atom being in the plane of the paper.
Shapes of more complicated molecules
Most molecules don’t consist of one central atom with a few other atoms bonded to it. In a larger molecule, we need to consider each atom that has more than one other atom bonded to it. While we can’t specify the overall shape of the molecule we can specify the shape around at each of these atoms by considering lone pairs and bonding pairs in the usual way.
For example, consider the aldehyde, ethanal: CH3CHO
We could draw a dot-cross diagram for this molecule, or just consider the valence shell electrons of the bonded atoms, to determine that the C of the CH3 has four bonding pairs and no lone pairs. The shape around this carbon is therefore tetrahedral and the bond angles will be 109.5°. The other carbon, in the CHO group, has three bonding pairs (one of which is a double bond, but for shapes we treat these the same as single bonds) and no lone pairs. The shape around this carbon is therefore trigonal planar and the bond angles will be 120°.