Hypervalent Molecules

Introduction

PCl₅, SF₆ and XeF₄ are routinely drawn with more than eight electrons around the central atom. Their shapes are real and VSEPR handles them usefully. The older explanation that phosphorus or sulfur must promote electrons into low-lying d orbitals is not required by modern bonding evidence. A Lewis drawing is a bookkeeping model, while actual bonding can be polar and delocalized.

Core explanation

Count domains first. PCl₅ has five P–Cl bonding regions and no central lone pair in its simple Lewis drawing, giving a trigonal bipyramid: three equatorial chlorine atoms about 120° apart and two axial chlorine atoms above and below. Axial–equatorial angles are about 90°. SF₆ has six S–F bonds, giving an octahedral arrangement with adjacent F–S–F angles of about 90° and opposite sites 180° apart. XeF₄ has four Xe–F bonds plus two xenon lone pairs; six domains form an octahedral framework, with the lone pairs opposite and the four fluorines square planar.

In a Lewis drawing, phosphorus in PCl₅ appears surrounded by ten bonding electrons and sulfur in SF₆ by twelve. That does not mean ten or twelve electrons are all localized in a compact valence shell around the central atom. Fluorine and chlorine draw electron density toward themselves; bonding can be described with ionic contributions and delocalized molecular orbitals. For axial bonds in a trigonal-bipyramidal species, three-centre, four-electron descriptions can be useful. The exact orbital treatment depends on the molecule, but it does not require large 3d participation by P or S.

XeF₄ illustrates another point: a noble gas can form compounds with very electronegative fluorine under suitable conditions. Xenon's eight valence electrons and four fluorine atoms give 8 + 4(7) = 36 valence electrons. Four Xe–F single bonds use eight, completing terminal fluorine octets uses 24 more, and four electrons remain as two Xe lone pairs. Placing those lone pairs opposite each other gives square-planar F positions and cancels bond dipoles in the ideal molecule.

“Hypervalent” is a description of a formal electron-counting representation, not a mechanism by itself. The same species may be modeled through resonance, polar bonds and multicentre interactions. Old textbook hybrid labels such as sp³d or sp³d² can name a shape mnemonic, but presenting them as proof that d orbitals dominate the bonds is misleading. East Tennessee State University's advanced inorganic teaching discussion at https://chem.libretexts.org/Courses/East Tennessee State University/CHEM 4110%3A Advanced Inorganic Chemistry/03%3A Simple Bonding Theories/3.01%3A Lewis Electron-Dot Diagrams/3.1.03%3A Breaking the Octet rule in Hypervalent Atoms notes that realistic calculations find P and S 3d orbitals too high in energy for substantial bonding in PF₅ and SF₆. OpenStax gives XeF₄ geometry at https://openstax.org/books/chemistry-2e/pages/7-6-molecular-structure-and-polarity.

Step-by-step reasoning

1. Count all valence electrons and build a formal Lewis structure. 2. Count bonds and central lone pairs as electron domains. 3. Use VSEPR to predict the geometry and distinguish axial from equatorial positions. 4. Treat the expanded-octet drawing as electron bookkeeping, not literal central electron localization. 5. Explain bonding with polar and delocalized contributions when the simple model is insufficient.

Visual explanation

Draw PCl₅ as an equatorial triangle crossed by a vertical axial line, SF₆ as four equatorial points plus top and bottom points, and XeF₄ as only the square's four atom points with lone pairs above and below. Label shapes rather than adding d-orbital lobes.

Real-world analogy

A city map can draw five roads meeting at one junction without claiming every vehicle sits at the junction. Lewis lines meeting at P in PCl₅ show connectivity and electron accounting; the actual electron density spreads along polar bonds and over several atoms.

Real-world example

SF₆ has a very symmetric octahedral molecular shape and has been used as an insulating gas in electrical equipment. Its high greenhouse impact motivates alternatives in some applications. The useful property is tied to molecular stability and dielectric behaviour, not evidence for a literal d-orbital hybrid picture.

Why?

Why do XeF₄'s lone pairs sit opposite? In an octahedral domain arrangement, opposite positions maximize their separation, reducing lone-pair repulsion. The remaining four fluorine positions form a square plane, which also allows their bond dipoles to cancel ideally.

Common misconception

“PCl₅ exists, so phosphorus must use 3d orbitals for five ordinary two-centre bonds” overinterprets a Lewis diagram. Modern orbital calculations and alternative multicentre descriptions show that substantial 3d hybridization is unnecessary.

Worked example

For XeF₄, total valence electrons are 36. Four Xe–F bonds and completed F octets leave two lone pairs on xenon. Six domains give octahedral electron-domain geometry. With lone pairs trans to one another, the molecular shape is square planar. The example predicts geometry while remaining neutral about a simplistic expanded-shell mechanism.

Quick check

1. What are the molecular geometries of PCl₅ and SF₆? Answer: Trigonal bipyramidal for PCl₅ and octahedral for SF₆.

Exam focus

Use the standard VSEPR shape names and count domains correctly. If asked about bonding, distinguish a formal expanded octet from modern delocalized descriptions. Do not claim d-orbital hybridization is experimentally required merely because a diagram has five or six bond lines.

Advanced insight

Electron localization in hypervalent molecules can be analysed with molecular orbitals and natural bond orbitals; these often assign significant ionic character to bonds with electronegative ligands. Different localized descriptions may represent the same overall electron density, so a single hybridization label need not be a unique physical fact.

Summary

PCl₅ is trigonal bipyramidal, SF₆ octahedral and XeF₄ square planar. VSEPR and Lewis counting predict these geometries, while modern bonding accounts use polar and delocalized interactions rather than requiring major central d-orbital involvement.

Practice questions

1. How many electron domains surround xenon in XeF₄? Answer: Six: four bonding domains and two lone-pair domains. 2. Which positions are axial in PCl₅? Answer: The two positions above and below the equatorial three-atom plane. 3. Why is an expanded Lewis octet not proof of central d-orbital bonding? Answer: Lewis lines are bookkeeping; real electron density can be polar and delocalized, and calculations show little need for major P or S 3d contribution.