Problem 43 - A Complete Resource Book in Chemistry for JEE Main

Question

Both \(\mathrm{BF}_{3}\) and \(\mathrm{NF}_{3}\) are covalent but \(\mathrm{BF}_{3}\) molecule is non-polar while \(\mathrm{NF}_{3}\) is polar because (a) atomic size of boron is smaller than nitrogen (b) \(\mathrm{BF}_{3}\) is planar but \(\mathrm{NF}_{3}\) is pyramidal (c) boron is a metal while nitrogen is gas (d) BF bond has no dipole moment while NF bond has dipole

Step-by-Step Solution

Answer

(b) \( \mathrm{BF}_{3} \) is planar but \( \mathrm{NF}_{3} \) is pyramidal.

Step 1: Identify Bond Types

Both \( \mathrm{BF}_{3} \) and \( \mathrm{NF}_{3} \) are made through covalent bonding, meaning that electrons are shared between atoms. This makes both molecules covalent in nature.

Step 2: Molecular Geometry Analysis

\( \mathrm{BF}_{3} \) is a planar and symmetrical molecule with a trigonal planar geometry. The bond angles are each \( 120^{\circ} \), allowing for balance of any dipole moments. In contrast, \( \mathrm{NF}_{3} \) has a trigonal pyramidal shape owing to the lone pair on nitrogen, causing an uneven distribution of charges.

Step 3: Dipole Moment and Polarity

In \( \mathrm{BF}_{3} \), the molecular symmetry cancels out any bond dipoles, leading to a net dipole moment of zero. For \( \mathrm{NF}_{3} \), the presence of a lone pair on the nitrogen atom makes the molecule asymmetrical, resulting in a net dipole moment.

Step 4: Conclusion on Polar Nature

Thus, while both molecules are covalent, the geometry of \( \mathrm{BF}_{3} \) leads to it being non-polar, whereas the shape and bond polarity of \( \mathrm{NF}_{3} \) contribute to its nature as a polar molecule.

Key Concepts

Covalent Bonding

Covalent bonding is a fundamental type of chemical bond where two atoms share electrons. This type of bond is prevalent in organic compounds and is crucial for forming stable molecules. Typically occurring between nonmetals, covalent bonds are characterized by the sharing of electron pairs.

In molecules like \( \mathrm{BF}_{3} \) and \( \mathrm{NF}_{3} \), the atoms involved distribute electrons in such a way that each atom attains a more stable electronic configuration.

• In \( \mathrm{BF}_{3} \), boron forms bonds with three fluorine atoms by sharing electrons, resulting in a highly symmetrical distribution.
• Similarly, in \( \mathrm{NF}_{3} \), nitrogen shares electrons with three fluorine atoms. However, a lone pair of electrons on nitrogen affects the overall molecular shape and polarity.

Understanding covalent bonding lays the groundwork for exploring more complex interactions like molecular geometry and polarity.

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms in a molecule. This geometry dictates many of a molecule’s physical and chemical properties. It is directly influenced by the number and type of bonds, as well as the presence of any lone electron pairs.

For example, \( \mathrm{BF}_{3} \) displays a trigonal planar geometry. Its symmetric shape arises because the three boron-fluorine bonds spread out evenly at \( 120^{\circ} \) angles, resulting in a flat, triangular shape.In contrast, \( \mathrm{NF}_{3} \) has a trigonal pyramidal shape. This difference is due to the lone pair of electrons located on the nitrogen atom, which pushes the bonding pairs closer together, disrupting the planarity seen in \( \mathrm{BF}_{3} \).

Hence, molecular geometry not only describes the shape of a molecule but also helps predict reactivity and interactions with other molecules.

Dipole Moment

The dipole moment is a measure of the separation of positive and negative charges in a molecule. It’s a vector quantity, meaning it has both magnitude and direction, and it is a key factor in determining molecular polarity.

In \( \mathrm{BF}_{3} \), although each \( \text{B-F} \) bond possesses a dipole due to the difference in electronegativity between boron and fluorine, the molecule's symmetric planar shape ensures that these dipoles cancel out. This results in a net dipole moment of zero, classifying \( \mathrm{BF}_{3} \) as non-polar.On the flip side, \( \mathrm{NF}_{3} \) experiences a net dipole moment. The lone pair on nitrogen pushes the \( \text{N-F} \) bonds away, making them not cancel perfectly. Consequently, \( \mathrm{NF}_{3} \) becomes a polar molecule.

Understanding the dipole moment helps recognize why some molecules interact more strongly with electric fields or other polar substances.

Trigonometric Shapes

In chemistry, trigonometric shapes illustrate the spatial arrangement of atoms around a central atom in a molecule, largely determined by bonding and lone pairs of electrons. Recognizing these shapes is integral in predicting and explaining molecular behavior.

\( \mathrm{BF}_{3} \) is a classic example of a molecule with a trigonal planar shape. This configuration can be visualized as a flat triangle, with the central atom at the center and three identical atoms at its vertices.For \( \mathrm{NF}_{3} \), the presence of a lone electron pair changes the shape to trigonal pyramidal. This is akin to a three-legged stool with the nitrogen atom at the top and each leg representing a \( \text{N-F} \) bond.

• Trigonal planar geometry involves 120-degree angles, leading to a symmetrical and balanced shape.
• Trigonal pyramidal geometry involves slightly smaller angles due to the lone pair, making the structure more asymmetrical.

The distinct trigonometric shapes are pivotal in determining dipole moments and subsequent polarity of molecules.