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Odd-electron molecules, called radicals, contain an odd number of valence electrons and therefore have at least one unpaired electron. Since they possess an odd total, not every atom can achieve an octet. The superoxide anion exemplifies this: with 13 valence electrons, one oxygen atom has only seven electrons and cannot reach eight. These molecules are inherently unstable and highly reactive.

Boron and beryllium, from groups 13 and 2 respectively, naturally form molecules with fewer than eight valence electrons around the central atom. Beryllium typically has four electrons, while boron has six. These electron-deficient atoms are highly reactive and readily accept lone pairs from other molecules. For example, boron trifluoride reacts with ammonia because boron's incomplete octet attracts nitrogen's lone pair.

A coordinate covalent bond, also called a dative bond, forms when one atom donates both electrons of a bonding pair to another atom. This occurs when electron-deficient atoms like aluminum or boron encounter molecules with lone pairs. In aluminum chloride reacting with ammonia, nitrogen donates its lone pair to aluminum's incomplete octet, creating a coordinate covalent bond that stabilizes the complex.

Elements from the third row onward have access to empty d orbitals in their valence shell, allowing them to hold more than eight electrons—typically up to 12 or 14. Phosphorus, sulfur, and iodine exemplify this expanded octet capacity. In tetrachloroiodide anion, iodine accommodates 12 valence electrons. Second-period elements like carbon and oxygen lack d orbitals and cannot form hypervalent compounds.

Hypervalent molecules have a central atom with more than eight valence electrons in their Lewis structure. After assigning bonding pairs and satisfying octets for all outer atoms, extra electrons remain and are placed on the central atom. Phosphorus pentachloride and sulfur hexafluoride are common examples. These molecules form only from third-period and heavier elements that possess available d orbitals.

Second-period elements possess only s and p orbitals in their valence shell—one 2s and three 2p orbitals—providing a maximum of four orbital pairs. This limits them to holding only eight valence electrons total. Unlike third-period and heavier elements with accessible d orbitals, carbon and oxygen cannot expand their octets and therefore never form hypervalent compounds.

Begin by counting total valence electrons and drawing a skeleton structure. Distribute electrons as lone pairs on terminal atoms, prioritizing the most electronegative atoms for filled shells. Assign remaining electrons to the central atom. For odd-electron molecules, accept that not all atoms reach octets. For electron-deficient or hypervalent molecules, place extra electrons on the central atom to complete the structure accurately.