Problem 64 - A Complete Resource Book in Chemistry for JEE Main

Question

Which one of the following statements is true? (a) The dipole moment of \(\mathrm{NF}_{3}\) is more than \(\mathrm{NH}_{3}\) (b) The dipole moment of \(\mathrm{NF}_{3}\) is less than \(\mathrm{NH}_{3}\) (c) The dipole moment of \(\mathrm{NH}_{3}\) is zero (d) The dipole moment of \(\mathrm{NF}_{3}\) is equal to \(\mathrm{NH}_{3}\)

Step-by-Step Solution

Answer

(b) The dipole moment of \(\mathrm{NF}_{3}\) is less than \(\mathrm{NH}_{3}\).

Step 1: Understand the Concepts of Dipole Moment

The dipole moment is a measure of the separation of positive and negative charges in a molecule. It depends on the difference in electronegativity between atoms and the molecule's geometry.

Step 2: Examine the Geometry of NH3 and NF3

Both \(\mathrm{NH}_{3}\) and \(\mathrm{NF}_{3}\) are tetrahedral molecules with a lone pair on nitrogen, resulting in a trigonal pyramidal shape. However, the direction of the dipole moments depends on the highly electronegative fluorine atoms in \(\mathrm{NF}_{3}\) and less electronegative hydrogen atoms in \(\mathrm{NH}_{3}\).

Step 3: Calculate the Net Dipole Moment of NH3

In \(\mathrm{NH}_{3}\), the dipole moment vectors from each \(\mathrm{N-H}\) bond are pointed towards nitrogen due to its higher electronegativity compared to hydrogen, and these vectors add up to give a net dipole moment directed towards the nitrogen.

Step 4: Calculate the Net Dipole Moment of NF3

In \(\mathrm{NF}_{3}\), each \(\mathrm{N-F}\) bond's dipole moment vector points towards fluorine due to fluorine being more electronegative. However, the lone pair on nitrogen pushes the overall dipole vector slightly towards nitrogen, decreasing the net dipole moment compared to if the vectors pointed towards a highly polar negative atom.

Step 5: Compare Dipole Moments and Determine the Correct Answer

Given the understanding of how the dipole moments are calculated for these molecules, \(\mathrm{NH}_{3}\) has a higher net dipole moment compared to \(\mathrm{NF}_{3}\) because the vectors in \(\mathrm{NH}_{3}\) are more additive due to the lack of a counteracting highly electronegative element like fluorine in \(\mathrm{NF}_{3}\). Thus, option (b) is correct.

Key Concepts

Electronegativity

In chemistry, electronegativity refers to the ability of an atom in a molecule to attract electrons towards itself. It plays a crucial role in determining the nature of bonds between atoms.
Highly electronegative atoms tend to pull electron density towards themselves more forcefully. The difference in electronegativity between two bonded atoms can inform you whether a bond is polar or nonpolar. In polar bonds, electrons shared between atoms are unequally distributed, creating partial charges at different parts of the molecule.

• For example, in the \(\mathrm{NF}_{3}\) molecule, fluorine has a higher electronegativity compared to nitrogen, leading to dipole moments where the electron density is skewed towards fluorine.
• Contrastingly, in \(\mathrm{NH}_{3}\), nitrogen is more electronegative than hydrogen, causing the electron density to be concentrated towards the nitrogen atom.

Understanding electronegativity differences helps in predicting and explaining the physical and chemical behaviors of molecules, especially when analyzing dipole moments.

Molecular Geometry

Molecular geometry, also known as molecular structure, is the three-dimensional arrangement of atoms within a molecule. It significantly affects a molecule's physical properties and chemical reactivity. The geometry of a molecule influences how dipole moments from individual bonds add up.
For tetrahedral molecules with lone pairs, like \(\mathrm{NH}_{3}\) and \(\mathrm{NF}_{3}\), the shape is trigonal pyramidal. The positions of atoms and lone pairs impact where the dipole moments point and how they combine together.

• In \(\mathrm{NH}_{3}\), the dipole moments are additive due to the shape and arrangement of the more electronegative nitrogen, enhancing the overall dipole moment pointing towards the nitrogen.
• In \(\mathrm{NF}_{3}\), the electronegative fluorine atoms pull electron density away from nitrogen, and the dipole moments produce a smaller net dipole as their direction somewhat counteracts the influence of nitrogen's lone pair.

The understanding of molecular geometry is central to predicting how molecules behave in various environments and how they interact with other molecules.

Polar Molecules

Polar molecules have regions with slight electrical charges due to uneven distribution of electrons in their bonds. This molecular polarity is a result of both differences in electronegativity and the spatial arrangement of atoms (molecular geometry).
Such molecules have at least one polar bond, and they exhibit a net dipole moment if the polar bonds do not cancel each other out.

• For instance, \(\mathrm{NH}_{3}\) is a polar molecule since the \(\mathrm{N-H}\) bonds result in a significant overall dipole moment. They effectively focus electron density upwards towards the nitrogen atom due to its geometry.
• Meanwhile, \(\mathrm{NF}_{3}\) is also polar, yet its overall dipole moment is less pronounced due to the geometry and opposing influences of the electronegative fluorine atoms and the nitrogen's lone pair.

In essence, polar molecules like these are crucial in chemistry because they influence solubility, boiling and melting points, and interactions with other molecules. Being able to identify and rationalize molecular polarity forms a basis for understanding many chemical properties and reactions.