Octet rule

The bonding in carbon dioxide (CO₂): all atoms are surrounded by 8 electrons, fulfilling the octet rule.

The octet rule is a chemical rule of thumb that states that atoms of low (<20) atomic number 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.

The valence electrons can be counted using a Lewis electron dot diagram as shown at right for carbon dioxide. The electrons shared by the two atoms in a covalent bond are counted twice. In carbon dioxide each oxygen shares four electrons with the central carbon, and these four electrons are counted in both the carbon octet and the oxygen octet.

Explanation in quantum theory

The quantum theory of the atom explains the eight electrons as a closed shell with an s²p⁶ electron configuration. A closed-shell configuration is one in which low-lying energy levels are full and higher energy levels are empty. For example the neon atom ground state has a full n=2 shell (2s² 2p⁶) and an empty n=3 shell. According to the octet rule, the atoms immediately before and after neon in the periodic table (i.e. C, N, O, F, Na, Mg and Al), tend to attain a similar configuration by gaining, losing, or sharing electrons.

The argon atom has an analogous 3s² 3p⁶ 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 3s² 3p⁶ 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.

Example: sodium chloride

Ionic bonding is common between pairs of atoms, where one of the pair is a metal (such as sodium) and the second a non-metal (such as chlorine).

A chlorine atom has seven electrons in its outer electron shell, the first and second shells being filled with two and eight electrons respectively. The first electron affinity of chlorine (the energy release when chlorine gains an electron) is +328.8 kJ per mole of chlorine atoms. Adding a second electron to chlorine requires energy, energy which cannot be recovered by formation of a chemical bond. The result is that chlorine will very often form a compounds in which it has eight electrons in its outer shell (a complete octet).

A sodium atom has a single electron in its outermost electron shell, the first and second shells again being full with two and eight electrons respectively. To remove this outer electron requires only the first ionization energy) which is +495.8 kJ per mole of sodium atoms, a small amount of energy. By contrast, the second electron resides in the deeper second electron shell, and the second ionization energy required for its removal is much larger: +4562.4 kJ per mole. Thus sodium will, in most cases, form compounds in which it has lost a single electron and have a full outer shell of eight electrons, or octet.

The energy required to transfer an electron from a sodium atom to a chlorine atom (the difference of the 1st ionization energy of sodium and the electron affinity of chlorine) is small: +495.8 - 328.8 = +167 kJ mol^-1. This energy is easily offset by the lattice energy of sodium chloride: -787.3kJ mol^-1. This completes the explanation of the octet rule in this case.

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 exhibit more than four bonds, 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

1. ^ Abegg, R. (1904). "Die Valenz und das periodische System. Versuch einer Theorie der Molekularverbindungen (The valency and the periodical system - Attempt on a theory of molecular compound)". Zeitschrift für anorganische Chemie 39 (1): 330–380. doi:10.1002/zaac.19040390125. 
2. ^ Lewis, Gilbert N. (1916). "The Atom and the Molecule". Journal of the American Chemical Society 38 (4): 762–785. doi:10.1021/ja02261a002. 
3. ^ Langmuir, Irving (1919). "The Arrangement of Electrons in Atoms and Molecules". Journal of the American Chemical Society 41 (6): 868–934. doi:10.1021/ja02227a002.