The octet rule is a practical rule that explains the formation of the chemical bond of representative elements as a function of the electronic configuration of their valence shell.
According to this rule, the atoms combine with each other in such a way that each atom is surrounded by eight electrons in its valence shell (hence the word octet).
This rule was published simultaneously in 1916 by both Gilbert Lewis and Walther Kossel. It reflects the observation that noble gases are very stable and all, with the exception of helium, are characterized by having their valence shell filled with 8 electrons.
The theory states that atoms share or exchange electrons by forming chemical bonds , in order to acquire this full-shell electronic configuration.
The octet rule and Lewis structures are closely related. This is because the octet rule is one of the bases to understand the formation of the chemical bond, so it allows predicting in most cases, the Lewis structure of chemical substances.
In this sense, one of the essential steps to draw a Lewis structure, after placing single bonds between all the atoms that make up the molecule, is to distribute the remaining valence electrons among the most electronegative atoms to complete their octets.
In cases where octets remain unfilled, double or triple bonds are formed, always seeking to comply with the octet rule for as many atoms as possible. So this empirical rule is fundamental for the construction of Lewis structures.
Examples of compliance with the octet rule
The octet rule applies mainly to representative elements, that is, to those in the s and p blocks of the periodic table. This rule applies consistently to compounds containing carbon, nitrogen, oxygen, and halogens, and with some exceptions to the other elements of the p block. However, most of the transition metals do not comply.
Some examples of compounds where the octet rule holds for all atoms are:
Carbon dioxide (CO 2 )
As can be seen from the Lewis structure of carbon dioxide presented below, both the central carbon and both oxygen atoms comply with the octet rule.
Neopentane (C 5 H 12 )
In the vast majority of organic compounds , all atoms comply with the octet rule (except for hydrogen, which is surrounded by only 2 electrons). This includes alkanes like neopentane or 2,2-dimethylpropane:
In this case it can also be observed that each carbon atom shares its four electrons with 4 neighboring atoms, each of which contributes one of its electrons to form the 4 simple covalent bonds that hold them together.
For this reason, each carbon atom is surrounded by 8 valence electrons, complying with the octet rule.
Carbon monoxide (CO)
Carbon monoxide is another molecular compound in which oxygen and carbon share electrons forming a covalent bond to satisfy the octet rule.
In this case, carbon contributes two electrons and oxygen 4 to complete the 6 electrons required to form the triple covalent bond that joins them. Each of the two atoms has an extra pair of unshared electrons that make up each octet.
Nitrate anion (NO 3 – )
The octet rule also holds for ions. The nitrate ion has a central nitrogen atom surrounded by 3 oxygen atoms.
As can be seen in this case, all the atoms have their octet filled, two of the oxygens have a negative formal charge while nitrogen has a positive formal charge, which results in the net charge of -1 of the nitrate ion.
Sodium Chloride (NaCl)
The exchange of electrons to form ionic compounds is also a common example of applying the octet rule.
When sodium chloride is formed from chlorine and sodium, it starts with two atoms that do not comply with the octet rule, since sodium has over one electron and chlorine lacks one to complete its octet.
Then, sodium gives its electron to chlorine, leaving the cation surrounded by 8 electrons and in turn completing the octet to chlorine by forming the chloride ion.
Exceptions to the octet rule
Just as there are a large number of compounds in which all atoms meet the octet rule, there are also multiple examples of atoms that do not.
Some are surrounded by fewer than 8 electrons, making them electron-deficient species , while others are surrounded by more than eight electrons, in which case it is said to have an expanded octet or is called a hypervalent atom .
Examples of electron deficient species
Borano (BH 3 )
This compound is a typical example of an electron-deficient species. Neither boron nor its surrounding hydrogens possess enough electrons to satisfy the octet rule in the central atom. This allows boron to receive a pair of electrons from another atom in another molecule, turning borane into a Lewis acid.
Aluminum trichloride (AlCl 3 )
AlCl 3 is another example of a Lewis acid that owes its chemical behavior to not meeting the octet rule.
Examples of hypervalent species
Sulfur hexafluoride (SF 6 )
A typical example of an expanded octet is SF 6 , which is usually represented as a central sulfur with 6 simple covalent bonds with fluorine atoms. In this case, the sulfur is surrounded by 12 valence electrons instead of 8, thus violating the octet rule.
Despite this, models have been proposed in which sulfur is actually covalently bonded to 4 fluorines at a time, while with the other two, it forms ionic bonds. If so, it would meet the octet rule.
Phosphorous pentachloride (PCl 5 )
As in the case of sulfur, phosphorus can also form compounds with expanded octets, in this case surrounded by 10 electrons.