Bonding (hl)

The actual arrangement of the atoms around the central atom in a molecule.

In order to find this we need to consider the number of regions of electron density (electron pairs or unpaired single electrons in some cases) around the central atom. The electrons will repel as far apart as possible as they have the same charge.

The valence shell electron pair repulsion theory (VSEPRT) gives us a means of working out the shapes of molecules and ions.

1. Draw the Lewis structure to show all of the valence shell electrons.

2. Then count the number of regions of electron density - this can then be translated into the electronic shape.

3. Now consider the number of attached atoms and their orientation keeping the lone (non-bonding) pairs as far apart as possible without changing the electronic shape.

4. The resultant molecular arrangement gives the molecular shape. Note that it may be slightly distorted by repulsions between regions of electron density.

For a fuller account see VSEPRT

Molecules in three dimensions

Click here and when the viewing screen opens right click on the image to see the options.

note: the MDL Chime plug-in must previously be installed (download here)

	BF3	CH4	PCl5	SF6
	NH3	XeF4

Predicting molecular shapes

Summary of the possible geometries

No. of electron pairs around the central atom	electronic shape	no of attached atoms	molecular shape
2	linear	2	linear
3	trigonal planar	3	trigonal planar
4	tetrahedral	4	tetrahedral
4	tetrahedral	3	pyramidal
4	tetrahedral	2	angular or bent
5	trigonal bipyramidal	5	trigonal bipyramidal
5	trigonal bipyramidal	4	square pyramid
5	trigonal bipyramidal	3	trigonal planar
5	trigonal bipyramidal	2	angular
6	octahedral	6	octahedral
6	octahedral	5	umbrella
6	octahedral	4	square planar

 

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14.2 - Hybridisation

This model explains the tetrahedral geometry of carbon and other atoms.

The electron structure of carbon is 1s² 2s² 2p² suggesting that it should only be able to form two bonds (using the two singly occupied orbitals). However it is known to make four single bonds in many compounds and indeed never forms just two bonds. This can be explained by hybridisation - the mixing of atomic orbitals producing degenerate orbitals used for bonding.

• sp³ hybridisation occurs when the 2s and 2p orbitals merge to become sp³ orbitals (all of equal energy, length etc.).

• sp² is the same except only two of the p orbitals are hybridised, leaving one p orbital unchanged

• sp is the same except only one of the p orbitals is hybridised and two p orbitals are left unchanged

hybridisation	geometry	example	carbon in
sp3	tetrahedral	methane	CH4
sp2	trigonal planar	ethene	CH2=CH2
sp	linear	ethyne	CH=CH

 

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14.3 - Delocalisation of electrons

When a particular molecule can be represented as several different Lewis structures is is generally not actually any of these, but a hybrid (mixture) of all of them. This can be represented either by using delocalised electrons, or through resonance (where each possible structure is drawn and the actual state 'resonates' between them. The delocalisation of these pi electrons (which is effectively what happens) makes the molecule more stable (as evidenced by lower energy) and gives the bonds a shorter length than would be expected.

Examples:

• benzene

• O₃

• SO₄^2-