Problem 142 - A Complete Resource Book in Chemistry for JEE Main

Question

The electronegativity difference between \(\mathrm{N}\) and \(\mathrm{F}\) is greater than that between \(\mathrm{N}\) and \(\mathrm{H}\), yet the dipole moment of \(\mathrm{NH}_{3}(1.5 \mathrm{D})\) is larger than that of \(\mathrm{NF}_{3}\) \((0.2 \mathrm{D})\) This is because (a) in \(\mathrm{NH}_{3}\) as well as \(\mathrm{NF}_{3}\) the atomic dipole and bond dipole are in opposite directions (b) in \(\mathrm{NH}_{3}\) the atomic dipole and bond dipole are in the opposite directions whereas in \(\mathrm{NF}_{3}\) these are in the same direction (c) in \(\mathrm{NH}_{3}\) as well as in \(\mathrm{NF}_{3}\) the atomic dipole and bond dipole are in same direction (d) in \(\mathrm{NH}_{3}\) the atomic dipole and bond dipole and in the same direction whereas in \(\mathrm{NF}_{3}\) these are in opposite directions

Step-by-Step Solution

Answer

(d) In \( \mathrm{NH}_3 \), bond and lone pair dipoles align; in \( \mathrm{NF}_3 \), they oppose.

Step 1: Understand the Concept of Dipole Moment

A dipole moment is a measure of the charge separation in a molecule and depends on both the magnitude and direction of the individual bond dipoles. The overall dipole moment of a molecule is the vector sum of all its bond dipole moments.

Step 2: Determine Electronegativity Differences

Electronegativity is a measure of an atom's ability to attract and hold onto electrons. In the given case, Fluorine (F) is more electronegative than both Nitrogen (N) and Hydrogen (H), meaning it attracts electrons more strongly. Therefore, the N-F bond has a higher individual bond dipole than the N-H bond.

Step 3: Analyze Molecular Geometry of \( \mathrm{NH}_3 \) and \( \mathrm{NF}_3 \)

Both \( \mathrm{NH}_3 \) and \( \mathrm{NF}_3 \) have a trigonal pyramidal shape due to the presence of a lone pair on nitrogen, which influences the directionality of dipole moments within the molecules.

Step 4: Consider Direction of Bond and Lone Pair Dipoles in \( \mathrm{NH}_3 \)

In \( \mathrm{NH}_3 \), the bond dipoles from N-H bonds and the lone pair dipole are aligned in the same direction, enhancing the overall dipole moment of the molecule.

Step 5: Evaluate Dipole Alignment in \( \mathrm{NF}_3 \)

In \( \mathrm{NF}_3 \), while the N-F bonds create downward dipoles due to fluorine's high electronegativity, the lone pair dipole on nitrogen points upward, opposing the bond dipoles. This results in a reduced overall dipole moment.

Step 6: Conclusion Based on Dipole Alignment

Since in \( \mathrm{NH}_3 \), the bond dipoles and lone pair dipole align in the same direction, whereas in \( \mathrm{NF}_3 \), they oppose each other, \( \mathrm{NH}_3 \) has a larger dipole moment than \( \mathrm{NF}_3 \). Therefore, the correct answer is (d).

Key Concepts

Electronegativity

Electronegativity is the tendency of an atom to attract electrons towards itself. It is a fundamental property that impacts molecular interactions and shapes molecular geometry, thus influencing properties like dipole moments. In simple terms, electronegativity determines how "greedy" an atom is for electrons.
Fluorine is one of the most electronegative elements on the periodic table. This means it has a strong desire to pull electrons towards itself, more so than both nitrogen and hydrogen.

• When dealing with bonds like N-F, the electronegativity difference is significant, making the bond highly polar with a strong dipole.
• In comparison, the N-H bond is less polar because the electronegativity difference between nitrogen and hydrogen is smaller.

These differences in electronegativity play a crucial role in determining the dipole moments of molecules such as ammonia (NH_3) and nitrogen trifluoride (NF_3). Overall, understanding electronegativity helps predict molecular behavior, especially regarding their bonding and interaction with other molecules.

Molecular Geometry

The three-dimensional shape of a molecule, or its molecular geometry, greatly affects the molecule's properties, including its dipole moment. Both NH_3 and NF_3 have a trigonal pyramidal shape due to the presence of a lone pair of electrons on the nitrogen atom. This lone pair pushes the bonded atoms downward, giving the molecules their characteristic shape.
This particular geometry is crucial because it dictates the direction in which the molecule's bonds point and, subsequently, how the individual bond dipoles contribute to the overall dipole moment.

• In NH_3, all N-H bond dipoles and the lone pair point in the same general direction adding constructively to increase the molecule's overall dipole moment.
• In NF_3, although the N-F bond dipoles are strong, they are countered by the lone pair dipole on nitrogen pointing in the opposite direction, thus reducing the overall dipole moment.

Molecular geometry determines not just visual representation but also influences practical characteristics such as solubility, boiling point, and as seen here, dipole moment.

Bond Dipoles

Bond dipoles are vectors that represent the polarity of a bond within a molecule. Calculated as the product of the charge difference across the bond and the distance between the charges, each bond dipole has both magnitude and direction.
For example, in NH_3:

• The N-H bonds create dipoles where the electron density is pulled slightly towards the nitrogen due to difference in electronegativity, though not quite as strongly as N-F.
• These dipoles, combined with the lone pair direction, reinforce one another, resulting in a significant molecular dipole moment.

In the case of NF_3:

• The N-F bonds are more polar, but because their dipoles and the lone pair dipole point in opposing directions, these cancel out to some extent.

Understanding bond dipoles helps explain why molecules with similar geometries can have such different dipole moments.

Lone Pair Electron

The lone pair electron is a pair of valence electrons that are not shared with another atom, remaining unbonded, or as "lone pairs." These have a profound effect on both shape and dipole moments of molecules.
In molecular geometry like in NH_3 and NF_3, these lone pair electrons create a region of electron density that repels bonded atoms, affecting molecular shape and dipoles.

• In NH_3, the lone pair creates a dipole that aligns with the N-H bond dipoles, thus increasing the overall dipole moment.
• Conversely, in NF_3, the lone pair's effect is opposite to the N-F bond dipoles, resulting in a significantly reduced dipole moment.

Understanding the role of lone pair electrons aids in predicting and explaining a molecule's physical properties and its chemical behavior in reactions.