 ## Shapes of octahedral molecules with lone pairs

The situation with octahedral shaped molecule with lone pairs is more straight- forward than that with trigonal bipyramidal molecules. The reason for this should be fairly clear, all positions in an octahedral molecule are equivalent, the bond angles between all the atoms in the molecule are 900. Whereas in a trigonal bypramidal molecule we had to consider two different locations, the axial and equatorial positions this is not the case with octahedral shaped molecules. The image below shows the shapes of several octahedral molecules that contain lone pairs of electrons. Recall that we do not consider the lone of electrons when determing the final shape. Molecules with a single lone pair will always form square pyramidal molecules, since all positions in an octahedral molecule are equivalent, it does not matter where the lone pair goes.
Molecules with two lone pairs can reduce the repulsion between the lone pairs by placing them 1800 apart. This will result in the formation of a square planar molecule, as shown below: ### Shapes of ions

You may meet questions which involve working out the shapes of ion. Remember that a cation, that is a positively charged ion is formed when a species loses an electron and anions negatively charged ions are formed by the addition of electrons to a molecule.

For example what shape is the ClF4- ion?
As before use the VSEPR rules to work out the shape of this ion:

1. Chlorine is the central atom and it is in group 7. It has 7 valency electrons
2. Four fluorine are bonded to the central atom, each contributes 1 electron. So we have 4 electrons in total.
3. Charge of -1, this means it has gained an electron.
4. The total number of electrons in the valency shells is 12 electrons, dividing by 2 gives 6 electron pairs, so the shape will be based on a octahedral structure, there are 4 bonding pairs from the 4 Cl-F bonds, this means there are two lone pairs of electrons. ### Example 2: what shape is the ammonium ion, NH4+? As before using VSEPR rules:

1. Nitrogen is the central atom and it is in group 5. It has 5 valency electrons.
2. Four hydrogen are bonded to the central atom, each contributes 1 electron. So we have 4 electrons in total.
3. Charge of +1, this means it has lost an electron.
4. The total number of electrons in the valency shells is 8 electrons, dividing by 2 gives 4 electron pairs, so the shap will be based on a tetrahedral structure, there are 4 bonding pairs of electrons from the 4 N-H bonds, this means there are NO lone pairs in this molecule.

### The effect of multiple bonds on bond angles and shapes of molecules

The 3 molecules shown below all have shapes based on a tetrahedral arrangement around the central atom. However as we have seen that lone pairs of electrons require more space than bonding electron pairs and this results in the reduction of the bond angles between the bonding pairs as shown below. Molecules with multiple bonds, that is double or triple bonds between the atoms show a similar effect. Double bonds for example contain a higher electron density than single bonds and so will repel other bonding electron pairs in a manner similar to that of lone pairs. In the example below are shown 2 molecule which have a trigonal planar structure. In a trigonal planar molecule we might expect bond angles of 1200, however as shown below you can see that the double bond requires more space than the single bonds in trigonal planar molecules and the bond angles will be not be the expected 1200. ## Key Points

• To find the shape of a molecule it is necessary to check to see if it contains any lone or non-bonding pairs of electrons.
• Lone pairs of electrons take up more space than bonding pairs of electrons. This can have a dramatic effect on the expected bond angles in a molecule.
• In deciding on the final shape of a molecule the lone pairs are NOT taken into account. Only the bonding pairs of electrons connected to other atoms are considered when taking into account the final shape of a molecule.