The octet rule is a chemical rule of thumb that reflects the theory that main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. The rule is especially applicable to carbon, nitrogen, oxygen, and the halogens, although more generally the rule is applicable for the s-block and p-block of the periodic table. Other rules exist for other elements, such as the duplet rule for hydrogen and helium, and the 18-electron rule for transition metals. The valence electrons in molecules like carbon dioxide (CO2) can be visualized using a Lewis electron dot diagram. In covalent bonds, electrons shared between two atoms are counted toward the octet of both atoms. In carbon dioxide each oxygen shares four electrons with the central carbon, two (shown in red) from the oxygen itself and two (shown in black) from the carbon. All four of these electrons are counted in both the carbon octet and the oxygen octet, so that both atoms are considered to obey the octet rule.
Example: sodium chloride (NaCl)
The octet rule is simplest in the case of ionic bonding between two atoms, one a metal of low electronegativity and the other a nonmetal of high electronegativity. For example, sodium metal and chlorine gas combine to form sodium chloride, a crystal lattice composed of alternating sodium and chlorine nuclei. Electron density inside this lattice forms clumps at the atomic scale, as follows. An isolated chlorine atom (Cl) has two and eight electrons in its first and second electron shells, located near the nucleus. However, it has only seven electrons in the third and outermost electron shell. One additional electron would completely fill the outer electron shell with eight electrons, a situation the octet rule commends. Indeed, adding an electron to the produce the chloride ion (Cl−) releases 3.62 eV of energy. Conversely, another surplus electron cannot fit in the same shell, instead beginning the fourth electron shell around the nucleus. Thus the octet rule proscribes formation of a hypothetical Cl2− ion, and indeed the latter has only been observed as a plasma under extreme conditions. A sodium atom (Na) has a single electron in its outermost electron shell, the first and second shells again being full with two and eight electrons respectively. The octet rule favors removal of this outermost electron to form the Na+ ion, which has the exact same electron configuration as Cl−. Indeed, sodium is observed to transfer one electron to chlorine during the formation of sodium chloride, such that the resulting lattice is best considered as a periodic array of Na+ and Cl− ions. To remove the outermost Na electron and return to an "octet-approved" state requires a small amount of energy: 5.14 eV. This energy is provided from the 3.62 eV released during chloride formation, and the electrostatic attraction between positively-charged Na+ and negatively-charged Cl− ions, which releases a 8.12 eV lattice energy. By contrast, any further electrons removed from Na would reside in the deeper second electron shell, and produce an octet-violating Na2+ ion. Consequently, the second ionization energy required for the next removal is much larger – 47.28 eV – and the corresponding ion is only observed under extreme conditions.
History
In 1864, the English chemist John Newlands classified the sixty-two known elements into eight groups, based on their physical properties. 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 that 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. In 1916, 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. In 1919, Irving Langmuir refined these concepts further and renamed them the "cubical octet atom" and "octet theory". The "octet theory" evolved into what is now known as the "octet rule". Walther Kossel and Gilbert N. Lewis saw that noble gases did not have the tendency of taking part in chemical reactions under ordinary conditions. On the basis of this observation, they concluded that atoms of noble gases are stable and on the basis of this conclusion they proposed a theory of valency known as "electronic theory of valency" in 1916:
During the formation of a chemical bond, atoms combine together by gaining, losing or sharing electrons in such a way that they acquire nearest noble gas configuration.
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