The octet rule or eight-electron rule is a classic rule in chemistry . It says that the electron configuration of atoms of the main group elements from the second period of the periodic table in molecules is a maximum of eight outer electrons ( valence electrons ) or four pairs. The atoms therefore strive to adopt the noble gas configuration . The octet rule is therefore a special case of the more comprehensive noble gas rule .
Atoms that mostly behave according to the octet rule
The octet rule often only applies to the main group elements of the 2nd period. These include the elements carbon , nitrogen , oxygen and fluorine . In most of their compounds, these elements achieve the electron configuration of the noble gas neon . The carbon, nitrogen and fluorine atoms also have the neon electron configuration with eight valence electrons in their elementary state - as diamond or fullerene , as dinitrogen (N 2 ), tri-oxygen (O 3 , ozone ) and difluoromolecule (F 2 ). For all atoms mentioned, it is true that they can be surrounded by fewer than 8 electrons (e.g. as R 3 C + , carbocation ), but never by more than 8.
The octet rule applies to most stable connections of the above elements. But there are exceptions. There are significantly more exceptions for elements of the higher periods. Elements such as tin and lead , for example, also form cations in the divalent oxidation state with a lone pair of electrons ( relativistic effect ).
However, there are many molecules in which the electron octet is only formally exceeded ( octet expansion ). Typical examples are phosphorus pentafluoride (PF 5 ), sulfur hexafluoride (SF 6 ) or iodine heptafluoride (IF 7 ). Until now, efforts have often been made to find unoccupied d orbitals with higher energies for the binding electrons that go beyond the octet. However, more detailed quantum mechanical considerations show that the d orbitals should not play an essential role because of the enormous energy difference to the s and p valence orbitals. Alternative descriptions of these molecules use multicenter bonds or partially ionic formulations (e.g. PF 4 + F - , SF 4 2+ (F - ) 2 , IF 4 3+ (F - ) 3 ).
Molecules for which Lewis formulas conforming to the octet rule can be drawn up, but for which formulas with more than 4 dashes are often used, should not be counted as exceptions. Sulfuric acid and sulfur dioxide are typical examples here .
Hydrogen and the light cations Li + , Be 2+ and B 3+ do not meet the octet rule because they have too few electrons and the associated noble gas configuration (helium), which is achieved in compounds, only has two electrons. This is more of a formal classification, but they meet the noble gas rule .
Nitrogen and oxygen
Exceptions are, for example, the nitrogen oxides nitrogen monoxide NO, also known as nitrogen oxide, and nitrogen dioxide NO 2 . The molecules of these compounds are permanent radicals , so they have an odd number of electrons, which is fundamentally incompatible with the octet rule.
Another exception to the octet rule is the dioxygen molecule O 2 : Measurements show that it contains two unpaired electrons. The noble gas configuration, however, requires paired electrons. When potassium , rubidium and cesium reacts with air, the dioxygen molecule can transform into the hyperoxide ion O 2 - ; the hyperoxides KO 2 , RbO 2 and CsO 2 are formed . The hyperoxide ion has an odd number of electrons and therefore no octet either.
No validity for unstable intermediate stages
The rule applies above all to isolable connections. In many reactions, unstable but detectable intermediates occur that do not obey the octet rule, e.g. B. radicals such as the chlorine radical or carbocations in which the carbon has only six electrons, for example in the conversion of butyl chloride. Nitrenes and carbenes are also mentioned.
An analogous rule can be found for transition metal compounds: the 18-electron rule . However, this is only sufficiently well fulfilled - if at all - for complexes with predominantly covalently bound ligands, but even here there are countless exceptions.