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The octet rule is a chemistry rule of thumb that reflects the theory that main-group elements tend to chemical 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 .
The valence electrons in molecules like carbon dioxide (CO2) can be visualized using a Lewis structure. In Covalent bond, 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.
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 valence 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 electron-volt of energy. (2025). 9780130399137, Pearson Education Limited. ISBN 9780130399137 Per source, the enthalpy change for
is -349 kJ/mol. Unit conversion performed using WolframAlpha database, 13 April 2025. 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., p. 880. Source gives ionization energy of +495.8 kJ/mol. Unit conversion performed using WolframAlpha database, 13 April 2025. 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., p. 156. Source gives lattice energy of 783 kJ/mol. Unit conversion performed using WolframAlpha database, 13 April 2025. 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, p. 880. Source gives ionization energy of +4562 kJ/mol. Unit conversion performed using WolframAlpha database, 13 April 2025. – and the corresponding ion is only observed under extreme conditions.
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.See:
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 chemical 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 . 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 of are stable and on the basis of this conclusion they proposed a theory of valency known as "electronic theory of valency" in 1916:
The argon atom has an analogous 3s23p6 configuration. There is also an empty 3d level, but it is at considerably higher energy than 3s and 3p (unlike in the hydrogen atom), so that 3s23p6 is still considered a closed shell for chemical purposes. The atoms immediately before and after argon tend to attain this configuration in compounds. There are, however, some hypervalent molecules in which the 3d level may play a part in the bonding, although this is controversial (see below).
For helium there is no 1p level according to the quantum theory, so that 1s2 is a closed shell with no p electrons. The atoms before and after helium (H and Li) follow a duet rule and tend to have the same 1s2 configuration as helium.
Typically, octet rule violations occur in either low-dimensional coordination geometries or in radical species. Although it is commonly taught that hypervalent molecules violate the octet rule, ab initio calculations show that almost all known examples obey the octet rule. The molecules form many bond order through resonance (see below), but each resonance structure does obey the octet rule.
tend to adopt states in which the unpaired electron can delocalize through resonance. In such cases, the octet rule can be restored through the formalism of a 1- or 3-electron bond.
Species such as can be interpreted two different ways, depending on their spin state. Triplet carbenes are best thought of as two radicals localized on the same atom, and obey the octet rule in those radicals' shared spin-up orientation. Singlet carbenes tend to adopt a planar configuration, and are best thought of as obeying the planar sextet rule.
However other models describe the bonding using only s and p orbitals in agreement with the octet rule. A valence bond description of PF5 uses resonance between different PF4+ F− structures, so that each F is bonded by a covalent bond in four structures and an ionic bond in one structure. Each resonance structure has eight valence electrons on P.Housecroft C.E. and Sharpe A.G., Inorganic Chemistry, 2nd ed. (Pearson Education Ltd. 2005), p.390-1 A molecular orbital theory description considers the highest occupied molecular orbital to be a non-bonding orbital localized on the five fluorine atoms, in addition to four occupied bonding orbitals, so again there are only eight valence electrons on the phosphorus. The validity of the octet rule for hypervalent molecules is further supported by ab initio molecular orbital calculations, which show that the contribution of d functions to the bonding orbitals is small.Miessler D.L. and Tarr G.A., Inorganic Chemistry, 2nd ed. (Prentice-Hall 1999), p.48Magnusson, E., J.Am.Chem.Soc. (1990), v.112, p.7940-51 Hypercoordinate Molecules of Second-Row Elements: d Functions or d Orbitals?
Nevertheless, for historical reasons, structures implying more than eight electrons around elements like P, S, Se, or I are still common in textbooks and research articles. In spite of the unimportance of d shell expansion in chemical bonding, this practice allows structures to be shown without using a large number of formal charges or using partial bonds and is recommended by the IUPAC as a convenient formalism in preference to depictions that better reflect the bonding. On the other hand, showing more than eight electrons around Be, B, C, N, O, or F (or more than two around H, He, or Li) is considered an error by most authorities. In particular, instead of pentavalent N, tetravalent N+ is written (e. g. not H−O−N(=O)=O but H−O−N+(=O)−O−).
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