Problem 96 - A Complete Resource Book in Chemistry for JEE Main

Question

The correct increasing bond angle among \(\mathrm{BF}_{3}, \mathrm{PF}_{3}\) and \(\mathrm{ClF}_{3}\) follows the order (a) \(\mathrm{BF}_{3}<\mathrm{PF}_{3}<\mathrm{CIF}_{3}\) (b) \(\mathrm{PF}_{3}<\mathrm{BF}_{3}<\mathrm{CIF}_{3}\) (c) \(\mathrm{ClF}_{3}^{3}<\mathrm{PF}_{3}<\mathrm{BF}_{3}^{3}\) (d) \(\mathrm{BF}_{3}<\mathrm{PF}_{3}<\mathrm{CIF}_{3}\)

Step-by-Step Solution

Answer

The correct increasing bond angle order is \( \text{PF}_3 < \text{ClF}_3 < \text{BF}_3 \), so none of the options provided are correct.

Step 1: Understanding the Structures

Before determining the bond angles, we first need to understand the molecular geometries of \( \text{BF}_3 \), \( \text{PF}_3 \), and \( \text{ClF}_3 \). \( \text{BF}_3 \) is trigonal planar, \( \text{PF}_3 \) is pyramidal, and \( \text{ClF}_3 \) is T-shaped.

Step 2: Determining Bond Angles for BF3

\( \text{BF}_3 \) has a trigonal planar geometry. In such a geometry, the bond angle is 120° due to the equal repulsion between the three bonding pairs of electrons.

Step 3: Determining Bond Angles for PF3

\( \text{PF}_3 \) has a trigonal pyramidal geometry with a bond angle of approximately 107°. This is smaller than \( \text{BF}_3 \) due to the presence of a lone pair on phosphorus, which compresses the bond angles.

Step 4: Determining Bond Angles for ClF3

\( \text{ClF}_3 \) is T-shaped due to the presence of two lone pairs on chlorine, which reduces the bond angle further. The bond angles are approximately 90° or slightly less due to the lone pair-lone pair and lone pair-bonding pair repulsions.

Step 5: Arranging Bond Angles

Based on the determined angles, we arrange them: \( \text{PF}_3 \) (107°) < \( \text{ClF}_3 \) (~90°) < \( \text{BF}_3 \) (120°). Therefore, in increasing order of bond angles, the correct answer is \( \text{PF}_3 < \text{ClF}_3 < \text{BF}_3 \).

Key Concepts

Bond Angle

The bond angle is the angle formed between three atoms across at least two bonds. It is a critical feature of molecular geometry. For different molecules, the bond angle can inform us about how atoms are arranged in space. It is influenced by several factors: the number of electron pairs surrounding a central atom, the presence of lone pairs, and the type of atomic orbitals involved. The bond angle tells us how spread apart the atoms are in a molecule. This angle can help predict the molecule's physical and chemical properties, such as polarity, reactivity, and phase of matter.

• Electron Pair Repulsion: Atoms and lone pairs are negatively charged. They push each other away, affecting bond angles.
• Lone Pairs vs. Bond Pairs: Lone pairs repel more strongly than bond pairs, reducing the bond angle between bonding atoms.

Trigonal Planar

When a molecule has a trigonal planar shape, it means the atoms are arranged in a flat, triangular shape. This shape occurs when there are three atoms bonded to a central atom, with no lone pairs affecting the arrangement. A classic example of a trigonal planar molecule is \(\text{BF}_3\). In it, the boron atom is the central atom bonded symmetrically to three fluorine atoms.

• Bond Angle: The angle between bonds in a trigonal planar molecule is always 120°, as in \(\text{BF}_3\).
• Electron Counting: The central atom must have a total of three electron pairs for a perfect trigonal planar shape.

Trigonal Pyramidal

Trigonal pyramidal geometry resembles a pyramid, with a triangular base. This shape forms when there are three bonds and one lone pair on the central atom. A well-known molecule with this shape is \(\text{PF}_3\). Unlike the trigonal planar geometry, the lone pair influences the molecule’s symmetry.

• Bond Angle: In \(\text{PF}_3\), the bond angle is about 107°, smaller than a perfect tetrahedral angle due to the lone pair's stronger repulsion.
• Effect of Lone Pair: The lone pair occupies more space and pushes the bonding pairs closer together, compressing the bond angle.

T-shaped Geometry

T-shaped geometry occurs when a molecule has five regions of electron density but only three bonded atoms due to the presence of lone pairs. This geometry is less common compared to other molecular shapes. \(\text{ClF}_3\) is a typical example of a T-shaped molecule.

• Lone Pairs: There are two lone pairs that are on the central atom, chlorine, which necessitates this shape by compressing certain bond angles.
• Bond Angle: T-shaped molecules mainly have bond angles around 90°. In \(\text{ClF}_3\), they are slightly less than 90° due to the lone pairs pushing the fluorine atoms.
• Spatial Arrangement: Two lone pairs occupy equatorial positions, creating a T-shape for the three bonded atoms.