The octet rule was used in gcse chemistry to help us explain why elements reacted with each other, it was assumed the elements react in order to end up with full last electron shells, or noble gas electronic configurations. We assumed that elements end up with 8 electrons or a ns2np6 electronic configuration after they had reacted. However you will no doubt noticed in your studies in A-level chemistry that there are quite a few molecule which do not appear to follow the octet rule. It would probabaly be fair to say that its a rule and not a law. The octet rule tends to hold for most of the first 20 elements in the periodic table but there are many examples of elements which do not follow the octet rule, e.g.
The period 2 elements beryllium and boron can form compounds with fewer than 8 electrons around the central atom. Boron has an electronic configuration of 1s22s22p1 or just simply 2,3; based on this electron arrangement we might predict using the octet rule that boron would lose 3 electrons or gain 5 electrons in order to complete its octet of electrons; however it does neither but instead forms three covalent bonds to form a molecule with only 6 electrons in its valency shell. As an example consider the molecule boron trifluoride, BF3. A dot and cross diagram is shown below for this molecule. It is clear that the central boron atom has only 6 electrons in its valence shell. Aluminium another group III element also forms these trigonal planar molecules with only 6 electrons around the central atom. This would suggest that the octet rule does not apply to some group 3 elements. However these electron deficient molecules have empty electron sub-shells and are able to form dative covalent bonds which enable them to achieve an octet of electrons.
Beryllium, electronic configuration 1s22s2 also forms compounds which are electron deficient around the central atom. In beryllium chloride; shown below; the central beryllium atom has only 4 electrons in its electron outer shell.
There are many examples of molecules where the central atom has more than 8
electrons in its valency shell. These molecules
or ions with more than 8 electrons around the central atom are often referred to as being hypervalent.
These hypervalent elements are found mainly in period 3 and above in groups 3,4,5,6,7 of the periodic table
though the nobel gases krypton and xenon in group 0 will also form hypervalent compounds.
As an example consider:
Phosphorus is an element in group 5 and period 3
that can form hypervalent molecules e.g.
Phosphorus reacts with chlorine to form 2 chlorides: phosphorus
trichloride (PCl3) and phosphorus pentachloride (PCl5).
Phosphorus has an electronic configuration of [Ne]3s23p3 indicating that it has 5 electrons in its valence shell. It can form three covalent bonds with chlorine and so gain a share of 3 additional electrons to form PCl3 which will give the central phosphorus atom an octet of electrons and the same electronic configuration as the noble gas argon.
However the 3d electron sub-shell in phosphorus is very close in energy to the 3s and 3p sub-shells and one electron can be promoted from the 3s sub-shell to the empty 3d sub-shell. This will give phosphorus the new electronic configuration:
This new electronic configuration now gives phosphorus 5 unpaired electrons in its valency shell and allows it to form the hypervalent chloride PCl5 as shown below:
Sulfur has an electronic configuration of [Ne]3s23p4 indicating that it has 6 electrons in its valence shell. It can form 2 covalent bonds with another non-metal element and so gain a share of 2 additional electrons to form a compound which will give the central sulfur atom the same electronic configuration as the noble gas argon. Like phosphorus above the 3d electron sub-shell is very close in energy to the 3s and 3p sub-shells, so sulfur promotes 1 electron from the 3s sub-shell and also one electron from the 3p sub-shell into the empty 3d sub-shell. This will give the sulfur atom the new electronic configuration:
This new electron arrangement gives the sulfur atom 6 electrons in its valency shell and allows it to 6 covalent bonds, to form hypervalent molecules such as sulfur hexafluoride (SF6) as shown below:
The two examples given above are just 2 of the many hypervalent molecules that exist where the central atom in a molecule is a the non-metal element in period 3 and above. It is worth mentioning that in the two examples given above, PCl5 and SF6 that the corresponding molecules from the period above, namely NCl5 and OF4 do not exist and cannot be made. The main reason for this is due to the sizes of the central atoms. The period 3 elements P and S are large enough to fit 5 or 6 smaller F or Cl atoms around them, whereas the smaller period 2 nitrogen and oxygen atoms cannot fit as many atoms around them. We should also consider the availabilty of the d sub-shell for use in bond formation. In period 2 elements such as N and O only the 2s and 2p sub-shells are available for bonding.
Most of the molecules you will meet have an even number of electrons in their outer valency shell, however there is a rather curious group of molecules which contain an odd number of electrons in their valency shell, this obviously means that they cannot have an octet of electrons in their last shell as you cannot pair an odd number of electrons!
The following two examples show simple molecules with an odd number of electrons in their valency shell: