Question:

Incorrectly matched pair is:

Updated On: Apr 9, 2025

$XeO_3$ - pyramidal
$XeF_4$ - tetrahedral
$XeF_6$ - disorted octahedral
$XeOF_4$ - square pyramidal

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The Correct Option is B

Approach Solution - 1

Let's evaluate each of the compounds:
1. XeOF₄ – square pyramidal: Xenon oxyfluoride (XeOF₄) has a square pyramidal geometry. It has five bonding pairs of electrons (one oxygen and four fluorine atoms) and one lone pair of electrons on xenon, resulting in a square pyramidal shape.
2. XeO₃ – pyramidal: Xenon trioxide (XeO₃) has a trigonal pyramidal geometry, not a pyramidal shape as indicated. It has three bonding pairs of electrons and one lone pair on xenon, giving it a trigonal pyramidal geometry.
3. XeF₄ – tetrahedral: Xenon tetrafluoride (XeF₄) has a square planar geometry, not tetrahedral. It has four bonding pairs of electrons and two lone pairs on xenon, leading to a square planar structure.
4. XeF₆ – distorted octahedral: Xenon hexafluoride (XeF₆) indeed has a distorted octahedral geometry, due to the presence of lone pairs on xenon.

Thus, the incorrect pair is (B), where the geometry of XeO₃ is trigonal pyramidal, not pyramidal.

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Approach Solution -2

XeOF$_4$: Xenon oxyfluoride has a square pyramidal structure, which is correct.
XeO$_3$: Xenon trioxide has a pyramidal shape, which is correct.
XeF$_4$: Xenon tetrafluoride has a square planar geometry, not tetrahedral. Thus, this is the incorrect pair.
XeF$_6$: Xenon hexafluoride has a distorted octahedral shape, which is correct.

Hence, the incorrectly matched pair is XeF$_4$ - tetrahedral.

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Concepts Used:

VSEPR theory

The Valence Shell Electron Pair Repulsion Theory abbreviated as VSEPR theory is based on the premise that there is a repulsion between the pairs of valence electrons in all atoms, and the atoms will always tend to arrange themselves in a manner in which this electron pair repulsion is minimalized.

Postulates of Valence Bond Theory

The overlapping of two half-filled valence orbitals of two different atoms forms the covalent bond. The electron density between two linked atoms increases due to the overlapping. The molecule thus gains the property of stability.
More than one bond can be created if the atomic orbitals contain more than one unpaired electron. As per the valence bond theory, the paired electrons in the valence shell cannot participate in such bond formation.
A covalent bond has a specific direction. The region of overlapping atomic orbitals is also parallel to such a bond.
There are two types of covalent bonds based on the overlapping pattern: sigma bonds and pi bonds. The pi bond is formed by the overlapping of atomic orbitals sideways, whereas the sigma bond is created by the overlapping of atomic orbitals along the axis of the nucleus.

Types of Hybrid Orbitals:

Applications of Valence Bond Theory

The maximum overlap conditions which is given by the valence bond theory can explain the formation of covalent bonds in many compounds.
One of its most important applications is the variation in the length and strength of chemical bonds in H₂ and F₂ molecules can be explained by the difference in their overlapping orbitals.
The overlap of the 1s orbital of the hydrogen atom and the 2p orbital of the fluorine atom makes the covalent bond in an HF molecule which is explained by the valence bond theory.

Limitations of the VSEPR theory

The VSEPR model is not a theory. It does not explain or attempt to explain any observations or predictions. Rather, it is an algorithm that accurately predicts the structures of a large number of compounds.
VSEPR is simple and useful but does not work for all chemical species.
First, the idealized bond angles do not always match the measured values. For example, VSEPR predicts that and will have the same bond angles, but structural studies have shown the bonds in the two molecules are different by 12 degrees.
VSEPR also predicts that group-2 halides such as will be linear when they are actually bent. Quantum mechanics and atomic orbitals can give more sophisticated predictions when VSEPR is inadequate.