Problem 126 - A Complete Resource Book in Chemistry for JEE Main

Question

Match the following: List I List II 1\. CIF \(_{3}\) (i) planar 2\. \(\mathrm{NF}_{3}\) (ii) pyramidal 3\. \(\mathrm{BF}_{3}\) (iii) T-shaped 4\. \(\left[\mathrm{CrF}_{3}\right] \mathrm{x}\) The correct matching is: 1 2 3 4 (a) (-) (iii) (ii) (i) (b) (iii) (ii) (i) (-) (c) (iii) (i) (ii) (iv) (d) (iii) (ii) (i) (ii)

Step-by-Step Solution

Answer

The correct matching is (b) (iii), (ii), (i), (-).

Step 1: Analyze CIF_3

Chlorine trifluoride ( ClF_3 ) consists of a central Cl surrounded by three F atoms. It typically adopts a T-shaped molecular geometry due to the presence of two lone pairs on the chlorine atom. Therefore, it should be matched with (iii) T-shaped.

Step 2: Analyze NF_3

Nitrogen trifluoride ( NF_3 ) consists of a nitrogen atom bonded to three fluorine atoms. Due to the lone pair on the nitrogen atom, the molecule adopts a pyramidal geometry. Therefore, NF_3 should be matched with (ii) pyramidal.

Step 3: Analyze BF_3

Boron trifluoride ( BF_3 ) is composed of a central B atom connected to three F atoms. The absence of lone pairs on the boron allows the molecule to form a planar, trigonal geometry. Therefore, BF_3 should be matched with (i) planar.

Step 4: Match the options

Looking at the available options: - CIF_3 matches with (iii) T-shaped - NF_3 matches with (ii) pyramidal - BF_3 matches with (i) planar - [CrF_3]x is not specified further, but based on the given options and typical usage where it might not be directly matched, choose the option (b) which fits: (c) (iii) (ii) (i) (-)

Key Concepts

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. It plays a crucial role in determining many properties of a compound, such as polarity and reactivity. The shape of a molecule is primarily determined by the number of bond pairs and lone pairs around the central atom. This is where VSEPR (Valence Shell Electron Pair Repulsion) theory comes in handy. VSEPR theory helps predict the geometry of a molecule by considering that electron pairs will arrange themselves to minimize repulsion between them.

For instance, **Chlorine trifluoride (ClF extsubscript{3})** exhibits a T-shaped geometry with three bonded pairs and two lone pairs. The lone pairs slightly repel each other and the bonded pairs more than the bonded pairs repel each other, leading to the T-shape. On the other hand, **Boron trifluoride (BF extsubscript{3})** is planar because it has three bonded pairs and no lone pairs, forming a trigonal planar geometry. This shape allows for rotation in the plane, contributing to its linear and symmetric structure. **Nitrogen trifluoride (NF extsubscript{3})** adopts a pyramidal shape because of the three bonded pairs and one lone pair. The presence of the lone pair causes the bonded pairs to be pushed downward, resulting in a three-dimensional pyramid-like shape.

Chemical Bonding

Chemical bonding describes the process by which atoms attach and thereby stabilize each other to form molecules. There are three main types of chemical bonds: covalent, ionic, and metallic bonds. In the context of molecular geometry and structure, covalent bonding is particularly relevant.

In a covalent bond, atoms share electrons in order to achieve a more stable electron configuration. The type of atoms involved and the arrangement of electrons affect the bond's strength, length, and angle, which in turn affect the molecule's shape. For example, in **ClF extsubscript{3}, NF extsubscript{3},** and **BF extsubscript{3},** the central atom forms covalent bonds with the surrounding fluorine atoms. These bonds stabilize the structure and allow the molecules to exhibit their characteristic geometries such as T-shaped, pyramidal, and trigonal planar, respectively.
**T-shaped** configurations like in ClF extsubscript{3} reflect the significant influence of lone pairs on bonding, deviating from what would otherwise be a symmetrical arrangement. **Trigonally planar** structures, such as BF extsubscript{3}, show the effect of an atom having no lone pairs on a central atom, leading to equal bond angles and a flat shape. **Pyramidal** shapes in NF extsubscript{3} demonstrate how bonds and lone pairs compete for space, leading to a three-dimensional configuration.

Lewis Structures

Lewis structures are diagrams that represent the bonds between atoms in a molecule and any lone pairs of electrons. They are vital tools in predicting the molecular geometry using VSEPR theory.
These structures use dots to show electron pairs and lines to depict bonds. Drawing Lewis structures is a critical first step in understanding the shape and properties of a molecule. By looking at the Lewis structure, one can determine the number of bonds and lone pairs around a central atom, which influences the geometry.

For example, when drawing the Lewis structure of **ClF extsubscript{3},** one would show three single bonds between chlorine and fluorine atoms, along with two lone pairs of electrons on the chlorine atom. This predicts a T-shaped geometry. Similarly, **BF extsubscript{3}** would be drawn showing three boron-fluorine bonds and no lone pairs on boron, indicating a trigonal planar shape. **NF extsubscript{3},** meanwhile, would include three N-F bonds and one lone pair on nitrogen, consistent with a pyramidal shape.
By correctly interpreting Lewis structures, one can make accurate predictions about molecular geometry and gain insights into other fundamental properties such as polarity and bond angles.