Octet rule

The bonding in carbon dioxide (CO₂): all atoms are surrounded by 8 electrons, according to the octet rule. CO₂ is thus a stable molecule.

The octet rule is a simple chemical rule of thumb that states that atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas. The rule is applicable to the main-group elements, especially carbon, nitrogen, oxygen, and the halogens, but also to metals such as sodium or magnesium. In simple terms, molecules or ions tend to be most stable when the outermost electron shells of their constituent atoms contain eight electrons.

Explanation in quantum theory

The quantum theory is used to explain the valence electron energy shells. In short, an element's valence shell is full and most stable when it contains eight electrons, corresponding to an s²p⁶ electron configuration. This stability is the reason that the noble gases are so unreactive, for example neon with electron configuration 1s² 2s² 2p⁶. (Helium is an exception as explained below).

Note that a "full shell" means that there are the eight electrons in the valence shell when the next shell starts filling, even though higher subshells (d, f, etc.) have not been filled. There can be at most eight valence electrons in a ground-state atom because p subshells are always followed by the s subshell of the next shell. This means that once there are 8 valence electrons (when the p subshell is filled), the next additional electron goes into the next shell, which then becomes the valence shell.

A consequence of the octet rule is that atoms generally react by gaining, losing, or sharing electrons in order to achieve a complete octet of valence electrons. Reaction of atoms occurs primarily in two ways: ionically and covalently.

Some of the atoms for which the octet rule are most useful are:

• Carbon, C
• Oxygen, O
• The Halogens, F, Cl, Br, I

Exceptions

• The duet rule of the first shell - the noble gas helium has two electrons in its outer shell, which is very stable. (Since there is no 1p subshell, 1s is followed immediately by 2s, and thus shell 1 can only have at most 2 valence electrons). Hydrogen only needs one additional electron to attain this stable configuration, while lithium needs to lose one.
• Trivalent boron compounds such as BF₃ have only 6 electrons in the valence shell, as do some reactive species such as carbenes. These molecules often react so as to complete their octet: trivalent boron compounds are well known as Lewis acids which form a fourth bond with a Lewis base, and carbenes are even more reactive. Beryllium and aluminum can also have incomplete octets.
• Free radicals (e.g. nitric oxide) contain one or more atoms which have an odd number of electrons.
• Hypervalent molecules in which main group elements are bonded to more than four atoms, for example phosphorus pentachloride, PCl₅, and sulfur hexafluoride, SF₆. The bonding in such molecules has been controversial. One model considers that the P atom (in PCl₅) forms five true covalent bonds with the participation of a d orbital, in violation of the octet rule. However the currently preferred model uses three-center four-electron bonding and conforms to the octet rule. See also Hypervalent molecule#Bonding in hypervalent molecules.
• For transition metals, the 18-Electron rule replaces the octet rule, due to the importance of d orbitals in these atoms.

History

In the late 19th century it was known that coordination compounds (formerly called “molecular compounds”) were formed by the combination of atoms or molecules in such a manner that the valencies of the atoms involved apparently became satisfied. In 1893, Alfred Werner showed that the number of atoms or groups associated with a central atom (the “coordination number”) is often 4 or 6; other coordination numbers up to a maximum of 8 were known, but less frequent. In 1904 Richard Abegg was one of the first to extend the concept of coordination number to a concept of valence in which he distinguished atoms as electron donors or acceptors, leading to positive and negative valence states which greatly resemble the modern concept of oxidation states. Abegg noted that the difference between the maximum positive and negative valences of an element under his model is frequently eight.^[1] Gilbert N. Lewis referred to this insight as Abegg's rule and used it to help formulate his cubical atom model and the "rule of eight" which began to distinguish between valence and valence electrons.^[2] In 1919 Irving Langmuir refined these concepts further and renamed them the "cubical octet atom" and "octet theory".^[3] The "octet theory" evolved into what is now known as the "octet rule".

See also

• Lewis structure
• Electron counting
• 18-Electron rule

References