From the models of SF4, BrF3, and XEF4, deduce whether different atom arrangement

From the models of SF4, BrF3, and XEF4, deduce whether different atom arrangement, called geometrical isomers, are possible; if so, sketch them below. Indicate the preferred geometry for each case and suggest a reason for your choice. Indicate which structures have dipole moments and show their direction. Compound Electron Dipole Moment Preferred Reason Geometry Molecular (Polarity) geometry SF4 BrF3 XeF4 None 2. From the models of SF4, BrF3, and XeF4, deduce whether different atom arrangement, called geometrical isomers, are possible; if so, sketch them below. Indicate the preferred geometry for each case and suggest a reason for your choice. Indicate which structures have dipole moments and show their direction. Compound Electron Dipole SF4 BrF3 XeF4 N3™ 3. Using the Lewis structure predict the geometrical structures of the following ions and state the hybridization of the central atom. Ion Electron Geometry CO3²- NO3- BF4 Preferred Geometry Moment Molecular (Polarity) geometry Reason Molecular Geometry Formal Charge of Central Atom Central Atom Hybridization

The Correct Answer and Explanation is :

Here’s a breakdown of the molecular geometries, dipole moments, and hybridization of the given compounds and ions:

Part 1: SF₄, BrF₃, XeF₄ – Geometrical Isomers, Preferred Geometry, and Dipole Moments

1. SF₄ (Sulfur Tetrafluoride)

• Electron Geometry: Trigonal Bipyramidal
• Molecular Geometry: Seesaw
• Dipole Moment: Yes (Polar)
• Preferred Geometry: Seesaw shape due to the lone pair occupying an equatorial position to minimize electron repulsion.
• Reason: The lone pair distorts the ideal trigonal bipyramidal shape, creating an asymmetry that results in a net dipole moment.

1. BrF₃ (Bromine Trifluoride)

• Electron Geometry: Trigonal Bipyramidal
• Molecular Geometry: T-Shaped
• Dipole Moment: Yes (Polar)
• Preferred Geometry: T-Shaped structure with two lone pairs in equatorial positions.
• Reason: Lone pairs prefer equatorial positions in a trigonal bipyramidal arrangement to reduce electron pair repulsions. The asymmetry results in a net dipole moment.

1. XeF₄ (Xenon Tetrafluoride)

• Electron Geometry: Octahedral
• Molecular Geometry: Square Planar
• Dipole Moment: No (Nonpolar)
• Preferred Geometry: Square Planar with two lone pairs opposite to each other.
• Reason: The symmetrical distribution of lone pairs cancels out any dipole moment, making the molecule nonpolar.

Part 2: CO₃²⁻, NO₃⁻, BF₄⁻ – Geometry, Formal Charge, and Hybridization

1. CO₃²⁻ (Carbonate Ion)

• Electron Geometry: Trigonal Planar
• Molecular Geometry: Trigonal Planar
• Formal Charge of Central Atom: 0
• Hybridization: sp²

1. NO₃⁻ (Nitrate Ion)

• Electron Geometry: Trigonal Planar
• Molecular Geometry: Trigonal Planar
• Formal Charge of Central Atom: +1
• Hybridization: sp²

1. BF₄⁻ (Tetrafluoroborate Ion)

• Electron Geometry: Tetrahedral
• Molecular Geometry: Tetrahedral
• Formal Charge of Central Atom: 0
• Hybridization: sp³

Explanation (300 words)

The molecular geometries of SF₄, BrF₃, and XeF₄ are determined by VSEPR (Valence Shell Electron Pair Repulsion) theory. The presence of lone pairs affects their final shape. SF₄ adopts a seesaw shape as the lone pair occupies an equatorial position, minimizing electron repulsion. This asymmetry creates a net dipole moment, making it polar. BrF₃, with two lone pairs, takes a T-shaped geometry where the lone pairs occupy equatorial positions, resulting in a net dipole moment and polarity. XeF₄, despite being an octahedral electron geometry, adopts a square planar molecular shape due to two lone pairs positioned opposite each other, which cancels out dipole moments, making it nonpolar.

For the ions CO₃²⁻ and NO₃⁻, both have a trigonal planar structure with sp² hybridization, meaning that their bond angles are approximately 120°. In NO₃⁻, nitrogen carries a formal charge of +1 due to resonance, while in CO₃²⁻, the formal charge is evenly distributed. BF₄⁻ adopts a tetrahedral molecular geometry with sp³ hybridization, maintaining equal bond angles of 109.5°. The formal charge on boron is zero since all valence electrons are used in bonding.

This analysis highlights how lone pairs influence molecular geometry, dipole moments, and hybridization, confirming the structures of these molecules and ions.