14.1 - Shapes of molecules and ions
14.1.1:- State and predict the shape and bond angles using the VSEPR theory for 5- and 6-negative charge centres. The shape of the molecules/ions and bond angles if all pairs of electrons are shared, and the shape of the molecules/ions if one or more lone pairs surround the central atom, should be considered. Examples such as PCl5, SF6 and XeF4 can be used.
Electrons are negatively charged and, as such, repel each other. When electrons are in orbitals with the same energy (degenerate) these orbitals then orient themselves to be as far apart as possible so as to minimise repulsion. The orientations that are adopted will depend on the number of regions of electron density (orbitals) surrounding the nucleus.
If there are two regions of electron density attached to a nucleus then they will be on opposite sides of the nucleus i.e. they will be arranged 180 degrees to each other.
The following table summarises the shapes adopted by the regions of electron density
|
No of regions
|
shape adopted
|
angles between orbitals
|
|
2
|
linear |
180º
|
|
3
|
trigonal planar |
120º
|
|
4
|
tetrahedral |
109,5º
|
|
5
|
trigonal bipyramid |
120º and 90º
|
|
6
|
octahedral |
90º and 180º
|
It must be remembered that we are talking about the regions of electron density here and that this may not be reflected in the molecular shape. If the electrons are not being used to bond then they cannot be seen and the molecule must be described as if they were not there.
For example: water
The central oxygen atom has its electron arranged into four pairs in four distinct regions (the orbitals are sp3 hybridised)
These repel one another and adopt a tetrahedral arrangement.
However only two of the electron pairs are used in bonding and the other two pair are 'lone' (ie cannot be seen). The shape of the molecule is therefore 'angular' or 'bent'
Tetrahedral electronically but the molecular shape is angular
The valence shell electron pair repulsion theory now looks at the relative
strength of the repulsions between the lone pairs and the bonding pairs.
As the lone pairs are not drawn further away from the central atom by
another atomic nucleus then they exert a greater repulsion on each other
than the bonding pairs do on each other. Intermediate is the repulsion
felt between a bonding pair and a one pair.
Order of repulsion strength:
lone pair- lone pair >> lone pair - bonding pair >> bonding pair - bonding pair
This causes the tetrahedral electronic shape to distort and squeezes the
bonding pairs together. The bond angle then closes sightly from 109,5º
to 104,5º
Example: PCl5
In this molecule the phosphorus is the central atom and provides 5 electrons.
Each of the bonded chlorines provide one electron making 5 + 5=10 in total around the central atom.
Each bond has two electrons and there are five bonds using up 5 x 2 =10 electrons.
Hence there are no electron pairs left over and the central atom arranges the five regions of electron density into the ideal orientation - a trigonal pyramidal shape.
Phosphorus pentachloride
(click to see in 3D - if it doesn't work you have to download
chime)
Example XeF6
Xenon from group 0 has eight electrons in its outer shell and each fluorine provides one for the bond maing a total of 8 + 6=14 electrons.
It may be seen from the formula that the Xe-F bonds require a total of 6 x 2=12 electrons therefore there is also a lone pair to make up the 14 electrons.
Electronically the seven oritals orientate themselves to form a pentagonal bipyramidal shape with two different bond angles 90º and 72º. The lone pair will occupy the one of the orbitals at 90º to the plane of the central 5 orbitals but will distort them downwards according to the valence shell electron pair repulsion theory. The final result is a sort of umbrella shape.
Example XeF4
Xenon from group 0 has eight electrons in its outer shell and each fluorine provides one for the bond maing a total of 8 + 4 = 12 electrons.
It may be seen from the formula that the four Xe-F bonds require a total of 4 x 2 = 8 electrons therefore there are also two lone pairs to make up the 12 electrons.
Electronically the six oritals orientate themselves to form an octahedral shape with bond angles 90º . The lone pairs will occupy the orbitals as far apart as possible i.e. on oppositesides of the octahedral shape according to the valence shell electron pair repulsion theory.
The final result is a square planar arrangement of Xe-F bonds with lone pairs above and below.
Click to see in
3D - if it doesn't work you have to download chime
(The pink spheres represent the lone pairs)
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