Using the VSEPR theory
one can easily determine
what is repeling
what in a molecule
, however, determining the shape
of the mole
cule can be annoying at time
In order to help out with this problem here is a list that will help you determine the shape of a simple molecule given the number of electron pairs:
(Seeing as ACSII art isn't the best way to represent 3D molecules I'll draw them with ASCII and then describe them. Because of this, the diagrams will not be a true representation of the molecules' shape)
A molecule will be linear when it has one or two pairs of electrons in the valence level of the central atom and one or two bonding pairs attached to the central atom.
Eg: BeF2, H2, HCl
H - Cl or F - Be - F
Linear molecules look like a line
. They're straight
Triangular planar molecules look like a triangle with a dot in the middle of it. To make a triangular planar molecule you need three electron pairs in the valence level of the central atom and three bonding pairs attached to the central atom.
An example is BF3
The angle between each of the bonds is theoretically
If you can imagine a triangular pyramid with an atom in the middle you'll have a good idea of what a tetrahedral molecule looks like. Tetrahedral molecules have four electron pairs in the valence energy level of the central atom and four bonded pairs. The tetrahedral shape is one of carbon's favourites =).
CH4 and CF4 are good examples.
/ | \
H H H
If your central atom of the molecule has four pairs of valence electrons, but only three bonded pairs then it'll be pyramidal. It looks much the same as a tetrahedral, but without the atom in the middle. Instead of having the atom in the middle it is on the top.
Examples: NH3, PCl3
/ | \
H H H
What are those dots above the Nitrogen
? They're the lone
pair of electrons that are repelling the hydrogen
. I put them in so that you can see why the molecule isn't Triangular Planar.
If we again replace a bonded pair with a lone pair we'll get a different sort of shape. The V-shaped molecule is bent because the lone pairs of electrons are repelling the bonded pairs. The angle between the bonded pairs is 104.5 degrees.
Examples: H20, F2O
I kept this linear for last because it has lone pairs of electrons attached to the central atom. The other linear molecules don't. This molecule has one bonded pair of electrons and three lone pairs.
Examples: F2, Cl2
:F - F:
Those are the basic molecular structures
that you have to deal with in grade
in Western Australia
. I'm sure that there is alot more to it when it comes to organic chemistry
, but I don't have that knowledge. In spite
of that, you can always work it all out with valence shell electron pair repulsion theory
if you really need to =)
Oh, I almost forgot double
bonds. If a molecule has a double or a triple bond you treat it as if it were only one negative region
. Of course it'll have a stronger negative charge
, but it'll act
as only one region of negative charge.