The octet rule is a useful pattern, not a universal law. Some stable molecules cannot be represented with eight valence-shell electrons around every atom.

Three important classes of exceptions are:

• electron-deficient species, in which an atom has fewer than eight electrons around it;
• odd-electron species, whose total valence-electron count is odd;
• expanded-valence structures, in which some atoms from period 3 or below are represented with more than eight electrons around the central atom.

For example, $BF_3$ has only six electrons around boron in its simplest Lewis structure. Forcing a boron octet by introducing $B=F$ double bonds creates less favorable formal charges and is not the best elementary representation.

Nitric oxide, $NO$, has $5+6=11$ valence electrons. Because the total is odd, at least one electron must remain unpaired in any Lewis structure.

Species such as $PCl_5$ and $SF_6$ are commonly represented with five and six bonds around the central atom. Introductory Lewis bookkeeping describes these as expanded-valence structures.

When the octet rule conflicts with the actual electron count or produces implausible formal charges, preserve the correct total electron count and use the structure that best represents the species. More advanced bonding theories explain the limitations of the octet picture in greater physical detail.