Lewis Structures

Core Idea

Lewis structures are diagrams that show the arrangement of atoms and valence electrons in a molecule or polyatomic ion, using lines for bonding pairs and dots for lone pairs. The procedure: count total valence electrons (adjusting for ion charge), connect atoms with single bonds, distribute remaining electrons as lone pairs to satisfy octets, and convert lone pairs to multiple bonds if needed. Lewis structures are the foundation for predicting molecular geometry, polarity, and reactivity.

How It's Best Learned

Follow the step-by-step procedure systematically for dozens of molecules, including polyatomic ions. Cross-check by counting all valence electrons to ensure none are lost or gained. Practice molecules with expanded octets (like SO₃, XeF₄) and electron-deficient molecules (like BF₃).

Common Misconceptions

• The central atom is not always the most electronegative — it is usually the least electronegative atom (except hydrogen, which is always terminal).
• Lewis structures for ions must account for charge: add electrons for negative ions, subtract for positive ions.

Explainer

You already know from covalent bonding that atoms share electrons to fill their outer shells. Lewis structures are the tool that lets you see exactly how that sharing is arranged — which atoms are bonded to which, where the shared pairs sit, and where the unshared (lone) pairs reside. Every prediction about molecular shape, polarity, and reactivity starts from a correct Lewis structure, so mastering the drawing procedure is essential.

The procedure is systematic. First, count total valence electrons for all atoms in the molecule. For CO₂: carbon contributes 4, each oxygen contributes 6, giving 4 + 6 + 6 = 16 total. For polyatomic ions, add electrons for negative charges or subtract for positive charges (SO₄²⁻ gets 2 extra electrons). Second, identify the central atom — usually the least electronegative atom that is not hydrogen. Third, connect each outer atom to the central atom with a single bond (each bond uses 2 electrons). Fourth, distribute remaining electrons as lone pairs on the outer atoms to satisfy their octets. Finally, check the central atom: if it lacks an octet, convert lone pairs from adjacent atoms into double or triple bonds.

Applying this to CO₂: after placing single bonds (C−O−C uses 4 electrons), you have 12 electrons left. Distributing them as lone pairs on the oxygens gives each oxygen 3 lone pairs plus 1 bond = 8 electrons, but carbon has only 4 (two single bonds). Carbon needs more. Converting one lone pair from each oxygen into a bonding pair creates two double bonds: O=C=O. Now carbon has 8 electrons (two double bonds), each oxygen has 8 (two bonding pairs + two lone pairs), and all 16 valence electrons are accounted for.

Some molecules break the octet rule. Electron-deficient molecules like BF₃ have a central atom with fewer than 8 electrons — boron has only 6 in BF₃ and that is its most stable structure. Expanded octet molecules like PCl₅ or SF₆ have central atoms from period 3 or below that can accommodate more than 8 electrons using available d orbitals. When multiple valid Lewis structures can be drawn that differ only in the placement of electrons (not atoms), the molecule exhibits resonance — a concept you will explore next. The Lewis structure is not the final word on bonding, but it is always the first step.