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 \(\mathrm{NH}_{3}\)

Short answer

(b) The dipole moment of \(\mathrm{NF}_3\) is less than \(\mathrm{NH}_3\).

Step-by-step solution

Understanding Dipole Moment

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

Molecular Geometry of \(\mathrm{NH}_3\)

\(\mathrm{NH}_3\) (ammonia) has a trigonal pyramidal shape with nitrogen at the center and hydrogen atoms at the corners. The lone pair on nitrogen contributes to a net dipole moment pointing from the hydrogen atoms towards the nitrogen.

Molecular Geometry of \(\mathrm{NF}_3\)

\(\mathrm{NF}_3\) (nitrogen trifluoride) also has a trigonal pyramidal shape, but the electronegativity of fluorine is much greater than that of nitrogen, causing the bond dipoles to point from nitrogen towards the fluorine atoms. However, the overall dipole moment is in the opposite direction due to the influence of lone pairs.

Comparing Dipole Directions

In \(\mathrm{NH}_3\), the dipole moments due to the three \(\mathrm{N-H}\) bonds are reinforced by the lone pair, leading to a large net dipole moment. In \(\mathrm{NF}_3\), the **\(\mathrm{N-F}\)** bond dipole moments are partially canceled by the lone pair, leading to an overall smaller dipole moment than \(\mathrm{NH}_3\).

Conclusion and Selection

Comparing these factors, \(\mathrm{NH}_3\) has a larger dipole moment compared to \(\mathrm{NF}_3\) because the lone pair and bond dipoles in \(\mathrm{NF}_3\) are in conflicting directions, reducing the net dipole. Therefore, statement (b) is true: The dipole moment of \(\mathrm{NF}_3\) is less than \(\mathrm{NH}_3\).

Key concepts

Molecular Geometry

Molecular geometry is the three-dimensional arrangement of atoms within a molecule. This arrangement affects the molecule's physical and chemical properties, including its reactivity, polarity, and phase of matter. For molecules like ammonia (\(\mathrm{NH}_3\)) and nitrogen trifluoride (\(\mathrm{NF}_3\)), understanding their geometry is key to understanding their dipole moments.

The molecular geometry of \(\mathrm{NH}_3\) and \(\mathrm{NF}_3\) is termed as trigonal pyramidal. This means that the central atom, nitrogen, is bonded to three other atoms (either hydrogen or fluorine) and there is one lone pair of electrons on the nitrogen atom. This lone pair does not occupy the same atomic plane as the hydrogen or fluorine atoms, causing a pyramidal shape.

In both cases, the geometry affects how dipoles formed by individual atoms contribute to the overall dipole moment of the molecule. This is why even minor changes in geometry can lead to significant differences in molecular behavior.

Electronegativity

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. Different atoms have varying levels of electronegativity, contributing to the overall polarity of a molecule. In context, nitrogen and fluorine in \(\mathrm{NF}_3\) and nitrogen and hydrogen in \(\mathrm{NH}_3\) demonstrate this concept.

Fluorine is more electronegative than both nitrogen and hydrogen. In \(\mathrm{NF}_3\), this results in strong pull towards the fluorine atoms, creating bond dipoles that tend to point away from nitrogen. However, in \(\mathrm{NH}_3\), nitrogen is more electronegative than hydrogen, causing bond dipoles to direct towards the nitrogen atom, reinforcing the overall molecular dipole.

• This difference in electronegativity values among bonded atoms is why \(\mathrm{NH}_3\) ends up with a larger net dipole compared to \(\mathrm{NF}_3\).

Trigonal Pyramidal

The trigonal pyramidal molecular shape is a consequence of specific geometrical and electron configurations. For \(\mathrm{NH}_3\) and \(\mathrm{NF}_3\), this geometry arises because the nitrogen atom carries one lone pair and forms three bonds.

This results in a base of three atoms forming a triangular arrangement, with the nitrogen atom above the plane or pyramid peak. This shape affects the dipole moment because the molecular structure is not symmetrical. The presence of the lone pair also impacts the dipoles. Since it is not bonded, the lone pair can exert significant force in terms of molecular polarity, thus affecting measurements like dipole moment.

The shape also means that in \(\mathrm{NH}_3\), all bond angles are influenced by the lone pair, further increasing the net dipole moment compared to \(\mathrm{NF}_3\).

Lone Pair Effect

The lone pair of electrons on the nitrogen atom plays a crucial role in defining the dipole moment in both \(\mathrm{NH}_3\) and \(\mathrm{NF}_3\). Unlike bonded electrons, lone pairs are free to orient themselves in a way that contributes or contradicts the direction of the dipole moment.

In \(\mathrm{NH}_3\), the lone pair is positioned such that it reinforces the existing dipoles from the nitrogen-hydrogen bonds. This alignment results in a strong net dipole moment. Conversely, in \(\mathrm{NF}_3\), the situation is a bit different. The electronegativity of fluorine causes the dipoles to point towards the fluorine atoms, yet the lone pair pushes back the opposite direction.

• This opposition in \(\mathrm{NF}_3\) is why it ends up having a smaller dipole moment than \(\mathrm{NH}_3\).
• The discrepancy shows how significant the lone pair's effect is on the dipole moment of a molecule.