VSEPR/ Molecular Geometry Theory

By Gabrielle Stefou and Suzanne Monir

VSEPR stands for valence shell electron pair repulsion. It is a theory that is based upon the idea of electron domains, which are also referred to as electron clouds. These are regions of electron density surrounding the central atom in a molecule, which includes bonding electrons and non-bonding electrons (lone pairs). VSEPR theory states that, in a molecule, these electron clouds are trying to get as far away from each other as possible in three-dimensional space because these regions are negatively charged, and like charges repel. Because they want to get as far away from each other as possible, the terminal atoms will extend out at certain bond angles that optimize their distance from each other.

Some cool facts: VSEPR Theory was developed by scientists Ronald Nyholm  and Ronald Gillespie of McMaster university

There are two components of the VSEPR geometry of molecules: the electron domain geometry (basic structure), and the molecular geometry (the reported structure). The first component is like the “default” structure, but after you analyze other properties of the molecule like lone pairs, you will get an overall molecular geometry which may differ from the “default”. In this article, I will explain how to get from the Lewis Structure of the molecule to your final molecular geometry.

The first step in finding the VSEPR shape of your molecule is to draw out the Lewis structure. After you have done so, you must count the number of electron clouds that surround your central atom. This includes bonding and nonbonding electrons. For example, if you had drawn the lewis structure of carbon dioxide, you would count two electron clouds, because you have two double bonds surrounding the central carbon atom (Figure 1). If you had drawn the lewis structure of ClF5, you would count six electron clouds because you have five single bonds around the central chlorine and one lone pair (Figure 2).

Figure 2 – ClF5                                     [AX5E – Square Pyramidal]

Figure 1 – CO2            [ AX2 – Linear]

Default Structures

This number is very important because it will give you your predicted default structure. Here are the different default structures you can have based upon the number of electron clouds in your molecule:

1. One electron cloud = linear (bond angle: 180º)
2. Two electron clouds = linear (bond angle: 180º)
3. Three electron clouds = trigonal planar (bond angle: 120º)
4. Four electron clouds = tetrahedral (bond angle: 109.5º)
5. Five electron clouds = trigonal bipyramidal (bond angle: 90º and 180º)
6. Six electron clouds = octahedral (bond angle: 90º and 180º)
7. Seven electron clouds = pentagonal bipyramidal (bond angle: 72º and 90º)  **

** Note – this is less commonly discussed in secondary/first year studies  of molecular geometry

Now, as I have stated before, this just gives you your default geometry; you must further assess the lone pairs in your molecule to get the final geometry that you will state as your answer.

Molecular Geometry (inlcuding the Lone pairs)

To determine your final molecular geometry, you must count the number of lone pairs in your molecule, and consider what the default geometry of your molecule is. The reason why lone pairs make such a difference in your geometry is because the lone pairs take up more space, and thus compress the bond angles of the terminal atoms. One easy tip in determining your final structure is to set up your electron clouds in your default geometry (including your lone pairs), and then ignore the lone pairs to determine the final shape. Here are all the different combinations of molecular geometries with lone pairs:

Refer to A as the central atom ,

X as the atoms or ligands attached to the central atom;

E as the non-bonding electron pairs/lone pairs

** Note that the numbers for X and E will always add up to the number of binding sites in the default structure)

1. Linear (1 or 2 electron clouds) and 1 lone pair = linear

(3 binding sites or basic Trigonal Planar)

1. Trigonal Planar (3 electron clouds) withOUT lone pairs = triogonal planar (AX3)
2. Trigonal Planar (3 electron clouds) and 1 lone pair = bent  (AX2E)
3. Trigonal Planar (3 electron clouds) and 2 lone pairs = linear (AXE2)

(4 binding sites or basic Tetrahedral)

1. Tetrahedral (4 electron clouds) withOUT lone pairs = tetrahedral (AX4)
2. Tetrahedral (4 electron clouds) and 1 lone pair = trigonal pyramidal (AX3E)
3. Tetrahedral (4 electron clouds) and 2 lone pairs = bent (AX2E2)
4. Tetrahedral (4 electron clouds) and 3 lone pairs = linear (AXE3)

(5 binding sites or basic Trigonal Pyramidal)

1. Trigonal bipyramidal (5 electron clouds) withOUT lone pair = trigonal bipyramidal (AX5)
2. Trigonal bipyramidal (5 electron clouds) and 1 lone pair = see-saw (AX4E)
3. Trigonal bipyramidal (5 electron clouds) and 2 lone pairs = T shape (AX3E2)
4. Trigonal bipyramidal (5 electron clouds) and 3 lone pairs = Linear (AX2E3)

(6 binding sites or basic Octahedral)

1. Octahedral (6 electron clouds) withOUT lone pair = octahedral (AX6)
2. Octahedral (6 electron clouds) and 1 lone pair = square pyramidal (AX5E)
3. Octahedral (6 electron clouds) and 2 lone pairs = square planar (AX4E2)
4. Octahedral (6 electron clouds) and 3 lone pair = T-shaped (AX3E3)
5. Octahedral (6 electron clouds) and 4 lone pair = Linear (AX2E4)

The effects of lone pair(s) on the shape and the bond angles of the terminal atoms becomes clearer when you visualize where the bonds would lie with respect to the lone pair(s). That is why standard diagrams of the different geometries are used to aid in the explanation of these concepts. These diagrams look like:

Here, you can see that there are tetrahedral default shapes being represented by these diagrams. Here, the single solid lines are used to show terminal atoms in the same plane, dashed wedges are used to show terminal atoms in three-dimensional space extending backwards, and the solid wedge is shown to represent a terminal atom in three-dimensional space extending forwards. If you are ever asked to draw the shape of your molecule, you would use these diagrams to do so.

In essence, VSEPR theory is a very important concept that allows us to understand the shape of a molecule. If you take Biology this year, you will be sure to remember that size matters, shape matters, and orientation matters. The shape of a molecule will dictate its function, which impacts the way your body works at a metabolic level. Therefore, it is essential that we understand how to use VSEPR theory to understand the way we, and other organisms, function.

Have a look at our Chem-Is-Tree page to see how our classes had an interesting application to VSEPR

 

A summary of VSEPR theories:

Infographic credits:  Compoundchem

To view the VSEPR structures, use this University of Colarado’s Phet program

https://phet.colorado.edu/en/simulation/molecule-shapes