Limitations of Lewis Structures

The Third-Period Rule: Elements in the third period and below (e.g., Phosphorus, Sulfur, Chlorine) can accommodate more than eight electrons in their valence shell, sometimes holding 10 or 12 electrons.

Mechanism of Expansion: This expansion is possible because these atoms have access to relatively low-energy, empty d-orbitals within their valence shell ($n \ge 3$), allowing them to form more than four covalent bonds.

Environmental Factors: Expanded octets usually occur when a large central atom is bonded to small, highly electronegative terminal atoms like Fluorine or Oxygen.

Example Case: In a molecule like $SF_6$, the central Sulfur atom is surrounded by 12 electrons (6 bonding pairs), which is stable despite exceeding the octet.

Electron-Deficient Atoms: Some elements are stable with fewer than eight electrons. Hydrogen is the most common, completing its valence shell with only two electrons (a duet).

Group 2 and 13 Elements: Beryllium (Group 2) and Boron (Group 13) often form compounds where they have only four or six valence electrons, respectively.

Reactivity: Because these atoms lack a full octet, they are often very reactive toward molecules that can donate a pair of electrons to help them complete their octet.

Always Count First: Before drawing, calculate the total number of valence electrons. If the total is odd, you immediately know you are dealing with a free radical exception.

Check the Period: If the central atom is in Period 2, NEVER expand its octet. If you have extra electrons left over after satisfying terminal octets, they must go on the central atom only if it is Period 3 or higher.

Formal Charge Verification: Use formal charge to determine if an expanded octet is the most stable structure. Sometimes, expanding the octet reduces formal charges to zero, which is preferred.

Common Trap: Students often try to force an octet on Boron or Beryllium. Remember that $BF_3$ and $BeCl_2$ are stable with 6 and 4 electrons respectively.