Before reading this page make sure you know how to work out the basic shapes of molecules using the VSEPR model and it may help if you review this page which looks at how lone pairs affect tetrahedral shaped molecules: lone pairs in simple molecules. This page will look at how lone pairs of electrons affect the shapes of trigonal bipyramidal molecules.

Shapes of trigonal bipyramidal molecules with lone pairs

Trigonal bipyramidal (tbp) molecules often contain one or more lone pairs of electrons; this creates a bit of a dilemma since there are two different positions available in these molecules where the lone pair could be placed; that is the axial and the equatorial positions. So if a tbp molecule has a lone pair of electrons will it go into the axial or the equatorial position? Now recall that a lone pair of electrons requires more space than a bonding pair of electrons; so where in a trigonal bipyramidal molecule will the lone pair of electrons have the most space? In the axial or equatorial positions? The image below shows the lone pair in both the equatorial and the axial position in a tbp molecule.

Axial or equatorial? It's all a matter of repulsion

• The lone pair in the axial position will feel repulsion at 90° from the 3 equatorial bonding pairs of electrons and also from the other axial bonding pair. However this axial bonding pair is much further away and at an angle of 180° from the lone pair of electrons.


• The lone pair if placed in an equatorial position will feel repulsion from the 2 other equatorial bonding pairs at an angle of approximately 122° (the lone pair pushes the other equatorial bonding pairs further apart than the ideal 120°). The lone pair will also feel repulsion from the two axial bonding pairs of electrons which, like the equatorial bonding electron pairs, will be repelled and the usual angle of 90° will be increased due to this additional repulsion from the lone pair.

So we can say that in the axial position the lone pair of electrons will experience more repulsion since the three bonding pairs of electrons in the equatorial position are only 90° away whereas in the equatorial position there are two bonding pairs at 90° and the other equatorial bonding pairs are much further away. This means that in a trigonal bipyramidal molecule if there are lone pairs of electrons present they will occupy the equatorial positions.

Example 1- What shape is the sulfur tetrafluoride (SF₄) molecule?

Using the VSEPR model we have used before we can say:

1. Sulfur is the central atom in the molecule and being a group 6 non-metal in the periodic table it has 6 valence electrons in its outer shell.


2. Four fluorine atoms form 4 covalent bonds to the central sulfur atom and each fluorine atom will contribute 1 electron to the covalent bond with the central sulfur atom. So we have 4 electrons in total from the fluorine atoms.


3. The total number of electrons in the valence shells is therefore 10 electrons and dividing by 2 since each covalent bond involves the sharing of a pair of electrons gives 5 electron pairs; so the shape of the SF₄ molecule will be based on a trigonal bipyramidal shape. However the trigonal bipyramidal structure has 5 atoms bonded to the central atom but in SF₄ there are only 4 fluorine atoms around the central sulfur atom. This means that there is one lone pair or non-bonding pair of electrons in this molecule and it will occupy the equatorial position where there is more space for it.

Note: The bond angles shown above for the SF₄ molecule are for illustrative purposes only. The actual bond angles differ slightly from those quoted, due to additional electronic effects beyond A-level requirements. The key point is that the lone pair reduces the angles between all bonding pairs in the molecule; this is what the examiners expect you to know.

Example 2- What shape is the chlorine trifluoride (ClF₃) molecule?

1. Chlorine is the central atom in ClF₃ and it is found in Group 7 of the periodic table; so it has 7 valence electrons.
2. Three fluorine atoms are covalently bonded to the central chlorine atom with each fluorine atom contributing one electron to each covalent bond to the chlorine atom. So we have 3 electrons in total.
3. The total number of electrons in the valence shells is therefore 10 electrons; dividing by two gives 5 electron pairs; so the shape of the ClF₃ molecule will be based on a trigonal bipyramidal geometry. There will be 2 lone pairs of electrons (4 electrons in total) and 3 bonding pairs of electrons (6 electrons in total); one pair for each of the fluorine–chlorine bonds. The shape of this molecule is shown below; without the lone pairs of electrons present the molecule is simply described as T-shaped for obvious reasons!

However the presence of the two lone pairs of electrons in this molecule will distort the T-shaped molecule. The additional space requirements of the lone pairs will mean that the 180° angle formed by the F–Cl–F atoms will be compressed down to 175°, this is outlined in the image below:

Example 3- What shape is the xenon difluoride (XeF₂) molecule?

The noble gases are generally very unreactive; however the larger noble gas such as xenon will react with fluorine to form xenon difluoride (XeF₂). What shape will this molecule have? Well as before simply use the VSEPR rules we have used so far to work out its shape:

1. Xenon is the central atom in this molecule and it is found in Group 18 of the periodic table so it will have 8 valence electrons.


2. Two fluorine atoms are bonded to the central xenon atom and each fluorine atom will contribute 1 electron to the covalent bond with the central xenon atom. So we have 2 electrons in total from the two fluorine atoms.


3. The total number of electrons in the valence shells is therefore 10 electrons; dividing by 2 gives 5 electron pairs; so the shape of the XeF₂ molecule will be based on a trigonal bipyramidal geometry. There will be 3 lone pairs of electrons (6 electrons in total) and 2 bonding pairs of electrons; one pair for each of the xenon fluorine covalent bonds.
   
   The 3 lone pairs will occupy the equatorial positions with bond angles of 120°; the two fluorine atoms will be in the axial positions. The overall shape of the molecule (ignoring the lone pairs of electrons) is linear. The molecule is perfectly straight!

Self-check: Shape and Lone/bonding pair activity

Use the activity below to decide on the shape and bond angles present in various molecules. Decide if the bond angles are as you expected or if they are distorted by the presence of lone pairs in the molecule.

Key Points

• Trigonal bipyramidal molecules have two different positions around the central atom in a molecule — axial and equatorial.


• Lone pairs of electrons always occupy the equatorial positions because these positions provide more space and minimise repulsion.


• Lone pairs repel more strongly than bonding pairs, compressing bond angles and distorting the ideal trigonal bipyramidal geometry.


• The shape of a molecule is named using only the arrangement of the bonded atoms — lone pairs are ignored in the deciding on the final shape of a molecule.


• Examples to remember:
  • SF₄ → see-saw shape (1 lone pair)
  • ClF₃ → T-shaped (2 lone pairs)
  • XeF₂ → linear (3 lone pairs)
• The ideal trigonal bipyramidal bond angles (90°, 120°, and 180°) become distorted when lone pairs replace bonding pairs.


• At A-level, simply explain that lone pairs reduce bond angles; exact experimental values are beyond the scope of the A-level specification.

Check your understanding - Questions on shapes and lone pairs