The octet rule is used to describe the attraction of elements towards having, whenever possible, eight valence-shell electrons (4 electron pairs) in their outer shell. Because a full outer shell with eight electrons is relatively stable, many atoms lose or gain electrons to obtain an electron configuration like that of the nearest noble gas. Except for Helium (with a filled 1s shell), noble gases have eight electrons in their valence shells. The octet rule is used when drawing Lewis dot structures.

The components of table salt, sodium and chloride ions, illustrate the octet rule. Sodium (Na) with an electron configuration of 1s²2s²2p⁶3s¹ sheds its outermost 3s electron and, as a result, the Na⁺ ion has an electron configuration of 1s²2s²2p⁶. This is the same electron configuration as neon, a noble gas (i.e., highly stable and relatively nonreactive).

Chlorine (Cl), on the other hand, has an electron configuration of 1s²2s²2p⁶3s²3p⁵. Chlorine needs one electron to fill its outermost third shell with eight electrons. When chlorine takes on the electron shed by sodium then the Cl⁺ ion's electron configuration becomes 1s²2s²2p⁶3s²3p⁶. This is the same configuration as Argon, another noble gas.

The octet rule, however, does not accurately predict the electron configurations of all molecules and compounds. Not every nonmetal, nor metal, can form compounds that satisfies the octet rule. As a result, the octet rule must be used with caution when predicting the electron configurations of molecules and compounds. Some atoms violate the octet rule and surround themselves with more than four electron pairs.

In general, the octet rule works for representative metals (Groups IA, IIA) and nonmetals, but not for the transition, inner-transition or post-transition elements. These elements seek additional stability by having filled half-filled or filled orbitals d or f subshell orbitals.