Question

Mark out the incorrect match of shape for a given molecule/ion. (a) \(\mathrm{ICl}_{4}^{-}-\)square planar (b) \(\mathrm{NH}_{2}^{-}-\)Pyramidal (c) \(\mathrm{SF}_{4}-\) See \(-\) saw shape (d) \(\mathrm{XeOF}_{2}-\) Trigonal planar

Short answer

The incorrect match is (d) \\(\mathrm{XeOF}_{2}\rightarrow\) Trigonal planar.

Step-by-step solution

Identify the Molecular Geometry

Examine each molecule/ion and determine its electronic geometry based on its number of electron pairs and bonding pairs.

Assess ICl4- Shape

The molecule \(\mathrm{ICl}_{4}^{-}\) has iodine as the central atom with four chlorine atoms and two lone pairs, resulting in an octahedral electronic geometry. The structure becomes square planar due to lone pair repulsions, so this match is correct.

Assess NH2- Shape

The \(\mathrm{NH}_{2}^{-}\) ion has nitrogen as the central atom with two hydrogen atoms and one lone pair, leading to a bent or V-shaped molecular geometry, which makes the assertion of 'pyramidal' incorrect.

Assess SF4 Shape

The molecule \(\mathrm{SF}_{4}\) has sulfur as the central atom with four fluorine atoms and one lone pair, forming a trigonal bipyramidal electronic geometry. This configuration results in a see-saw shape, which is correct.

Assess XeOF2 Shape

The molecule \(\mathrm{XeOF}_{2}\) has xenon as the central atom with two fluorine atoms, one oxygen atom, and two lone pairs. Its electronic geometry is trigonal bipyramidal, while its molecular shape is T-shaped due to the two lone pairs. 'Trigonal planar' is not correct for this molecule.

Key concepts

Understanding Lone Pair Repulsions

Lone pair repulsions are a crucial concept in determining molecular geometry. Atoms in a molecule are surrounded by areas of electron density, which include both bonding pairs (the electrons involved in bonds between atoms) and lone pairs (the non-bonding electrons). The key principle here is that lone pairs require more space than bonding pairs. This is due to their position closer to the central atom, where they are less restricted by atomic bonding.

Because they occupy more space, lone pairs exert greater repulsive forces on other electron pairs. This repulsion alters the angles between bonds, leading to specific molecular shapes based on both the type and number of bonds in conjunction with lone pairs. For example,

• The ICl_{4}^{-}  ion is structured as a square planar shape due to the repulsions between its lone pairs bending the geometry from an octahedral matchup.
• SF_{4} shows another example; a lone pair on the sulfur atom pushes the neighboring fluorine atoms into a seesaw shape rather than a perfect trigonal bipyramidal arrangement.

Understanding this behavior helps in predicting and rationalizing the shapes that different molecules will adopt.

Exploring Octahedral Electronic Geometry

Octahedral geometry is a fascinating molecular structure in the world of electronic shapes. It is characterized by six groups of electrons surrounding a central atom. This configuration naturally adopts an arrangement where all six positions are equivalent, creating a perfect symmetry around the central atom.

For instance, in an ideal octahedral molecule like ICl_{4}^{-}, the presence of six electron domains leads to the six vertices of an octahedron all being occupied. However, when two of these positions are held by lone pairs, the molecular geometry shifts to a square planar shape. Such transformations occur because the lone pair-bond pair repulsion is greater than the bond pair-bond pair repulsion, leading to adjustments in the spatial configuration.

• This concept is not just theoretical; it has practical applications and implications in predicting how compounds will interact with each other.
• It also affects molecular stability and reactivity, making it a staple concept in advanced chemistry learning.

Understanding Trigonal Bipyramidal Geometry

Trigonal bipyramidal geometry stands out due to its arrangement of five electron groups around a central atom. This includes three atoms arranged in a plane (equatorial positions) and two atoms aligned axially (axial positions). This configuration is notably different from others due to its uneven distribution of electron pairs.

This geometry can accommodate different arrangements depending on the presence of lone pairs. For example, in SF_{4}, the presence of a lone pair distorts the shape away from perfect symmetry, resulting in a see-saw shape as opposed to a pure trigonal bipyramidal shape. Moreover, in the case of XeOF_{2}, two lone pairs cause a shift from trigonal bipyramidal to a T-shaped geometry.

• These changes in shape are defined by maximizing bond angle separation and minimizing lone pair repulsions.
• Trigonal bipyramidal geometry is integral to understanding the shapes and angles in complex molecules, often correlating to their physical and chemical properties.