The Octet Rule

Lewis’s major contribution to bonding theory was to recognize that atoms tend to lose, gain, or share electrons to reach a total of eight valence electrons, called an octet. This so-called octet rule  explains the stoichiometry of most compounds in the s and p blocks of the periodic table. We now know from quantum mechanics that the number eight corresponds to having one ns and three np valence orbitals filled, which together can accommodate a total of eight electrons.   We also know that the configuration ns²np⁶ is the one in a period which with the highest ionization energy and the lowest electron affinity.  This level is the most difficult to take a valence electron away from or add one to.  Atoms which can achieve an ns²np⁶ by sharing, borrowing or lending electrons to another atom which also achieves this configuration in the exchange will form a bond.

Remarkably, though, Lewis’s insight was made nearly a decade before Rutherford proposed the nuclear model of the atom and more than two before Schrodinger had explained the electronic structure of hydrogen.

For some time helium was treated as an exception to the octet rule.  Today we know that helium's 1s² electron configuration gives it a full n = 1 shell, and hydrogen, why gains its one electron to achieve the electron configuration of helium. We understand this as a consequence of only two electrons being able to fit in the n = 1 shell, in Lewis' time this was a mystery, something that was simply accepted.  It is the ability to understand the atomic orbital basis of ad hoc rules developed in the past that motivates our atoms first approach to chemistry.

Lewis dot symbols can also be used to represent the ions in ionic compounds. The reaction of cesium with fluorine, for example, to produce the ionic compound CsF can be written as follows:

No dots are shown on Cs⁺ in the product because cesium has lost its single valence electron to fluorine. The transfer of this electron produces the Cs⁺ ion, which has the valence electron configuration of Xe, and the F⁻ ion, which has a total of eight valence electrons (an octet) and the Ne electron configuration. This description is consistent with the statement   that among the main group elements, ions in simple binary ionic compounds generally have the electron configurations of the nearest noble gas. The charge of each ion is written in the product, and the anion and its electrons are enclosed in brackets. This notation emphasizes that the ions are associated electrostatically; no electrons are shared between the two elements.