This topic concerns the 3 dimensional shape of molecules. Molecular shape is

very important when chemical reactions are considered. This is especially true

in Biochemical reactions. This aspect of molecular shape will be dealt with in

the medicinal topic of the Organic Chemistry unit. For now we will limit the

discussion to simple molecules with relatively few atoms.

You are already familiar with the basic shape of methane and other

hydrocarbons.

The methane molecule has a tetrahedral geometry and the angle between each

of the bonds is 109.5 degrees

The question is - why does

a methane molecule adopt

this particular shape?

The basic answer to this question is due to the electrons in the covalent bonds

holding the atoms together.

Each covalent bond is a shared pair of electrons. As electrons are negatively

charged, the electrons in a bond will repel the electrons in other bonds.

This diagram shows the covalent bonds in methane.

Each bonded pair of electrons will repel the other

three bonded pairs of electrons.

The theory known as valence shell electron pair repulsion

(VSEPR) is used to explain why molecules adopt a particular

shape.

The basic principle behind the theory states:

Consider the molecule shown below. It contains atom X which has three

covalent bonds to the atoms of Y.

In this diagram two of the bonds angles are 90 degrees

and one is 180 degrees. This is an unfavourable geometry

as the repulsion of the bonded electrons is unequal –

there will be greater repulsion between the electron

pairs which are 90 degrees apart than between the

electron pairs which are 180 degrees apart.

In this diagram all the bond angles are all 120 degrees

and so the electron pairs are as far apart from each

other as possible. This will be the lowest energy situation

for this molecule as the repulsive forces are pushed as

far apart as possible.

In this example atom X has three bonded pairs of electrons – atom X is

bonded to three other atoms. This is not always the case. Molecules can have

more or less bonded pairs than the molecule shown (e.g. methane has 4 bonded

pairs). The number of pairs of electrons a molecule has will dictate the shape

adopted by the molecule. In addition some molecules have non – bonded pairs

of electrons called LONE PAIRS which also influence molecular shape.

Y

Y Y

Y

Y Y

When all the pairs of electrons are bonding pairs the repulsive force

generated by each pair will be very similar. This results in the FIVE basic

molecular geometries (shapes) shown below

Knowing the number of electron pairs around the central atom allows the

molecular geometry to be determined.

Example - BeCl2

2 Bonding Pairs

Example – BF3

3 Bonding Pairs

Example – CCl4

4 Bonding Pairs

Example – PCl5

5 Bonding Pairs

Example – SF6

6 Bonding Pairs

To determine the number of electrons pair around the central atom in a

molecule or ion we can use the following formula

If the particle is an ion the number of electrons lost or gained (the value of

the charge) must be added (- ion) or subtracted (+ ion) from the total number

of electrons before dividing by two.

1. Determine the number of electron pairs in the following molecules or ions.

a. H2O b. NH3 c. SF6 d. ICl4-

e. BCl3 f. NF3 g. SiCl4 h. PF5

i. PH4+ j. BeF2 k. IF5 l. SiCl62-

Another important factor in determining molecular geometry is the nature of

the electron pairs.

Consider the bonding in the ammonia molecule, NH3

The electronic configuration of nitrogen is

1s2 2s2 2p3

Nitrogen has three unpaired electrons in its valence shell . It is these electrons

which form the three bonds with the hydrogen atoms in ammonia. However,

nitrogen also has a pair of electrons in its valence shell – the 2 electrons in the

2s sub-shell. These are usually non – bonding electrons.

Ammonia has 4 pairs of electrons. Three are bonding pairs – hence three bonds,

and one lone pair {non bonding pair}.

The geometry of the ELECTRONS is TETRAHEDRAL but the geometry of the

ammonia molecule is a TRIGONAL PYRAMID.

Study the examples below which show how to determine the number of

bonding and non-bonding pairs in a molecule or ion.

This molecule has (8+4)/2 = 6 electron pairs.

This indicates the shape the electron pairs will be octahedral.

As the molecule has four F atoms connected to the central Xe atom it has four

bonding pairs of electrons and so it must also have two non- bonding pairs.

The shape the atoms adopt is based on the

octahedral geometry of the electron pairs but

as there are only four atoms on the central

xenon atom the shape of the molecule is

SQUARE PLANAR.

This molecule has (6+2)/2 = 4 electrons pairs

This indicates the shape the electron pairs will be tetrahedral.

As there are only two hydrogen atoms connected to the central oxygen it has

two bonding pairs of electrons and so it must have two non – bonding pairs of

electrons.

The shape the atoms adopt is based on the

tetrahedral geometry of the electron pairs but

as there are only two atoms on the central

oxygen atom the shape of the molecule is

described as NON-LINEAR or ANGULAR.

1. Go back to the question on page 5 and decide how many bonding and how

many non - bonding electron pairs each substance has.

In the exam you may be asked to draw a diagram of a molecule or ion which

clearly shows the shape of the molecule or the shape that the electron pairs

adopt. Your drawing must not look “FLAT” and must give an idea of the three

dimensional arrangement of the electron pairs or of the atoms.

The diagrams below show one way to do this.

Match the electron pair geometry of the molecules below to the shapes above.

a. BrF5 b. PCl4+ c AlBr3 d. XeF2 e. H3O+

Look back at page three to see the bond angles formed when all the electron

pairs around the central atom are bonding pairs.

If a molecule, or ion, has lone pairs of electrons on the central atom, the

shapes are slightly distorted away from the regular shapes. This is because of

the extra repulsion caused by the lone pairs.

As a result of the extra repulsion, bond angles tend to be slightly less as the

bonds are “squeezed” together.

The three molecules above show decrease in bond angle as the number of

NON-BONDED electron pairs exert a greater repulsive effect on the bonded

pairs. The same effect is seen in other molecules.

Remember that lone pairs of electrons cause bond angles to vary from normal.

A list of “learning outcomes” for the topic is shown below. When the topic is

complete you should review each learning outcome.

Your teacher will collect your completed notes, mark them,

and then decide if any revision work is necessary.

State that the shapes of simple molecules or ions can be explained by the

VESPR theory.

Be able to determine the number of bonding and the number of non-bonding

electron pairs around a central atom.

Be able to relate the shape of a molecule or ion to the number of electron pairs

around the central atom,

Be able to use draw perspective diagrams of the following shapes. Trigonal

planar, trigonal pyramid, tetrahedral, trigonal bipyramid, square planar and

octahedral.

State the bond angles the molecular shapes listed above.

State that lone pairs of electrons exert more repulsive force than bonded pairs

of electrons and that this results in molecules or ions with distorted shapes.

I have discussed the learning outcomes with my teacher.

My work has been marked by my teacher.

Teacher Comments.

Date. __________________________________

Pupil Signature. __________________________

Teacher Signature. _______________________

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