Problem 70 - General Chemistry: Principles and Modern Applications

Question

The structures of the \(\mathrm{NH}_{3}\) and \(\mathrm{NF}_{3}\) molecules are similar, yet the dipole moment for the \(\mathrm{NH}_{3}\) molecule is rather large (1.47 debye) and that of the NF \(_{3}\) molecule is rather small (0.24 debye). Provide an explanation for this difference in the dipole moments.

Step-by-Step Solution

Answer

The difference in the dipole moments of \(\mathrm{NH}_{3}\) and \(\mathrm{NF}_{3}\) can be explained by the directions of their bond dipoles. In \(\mathrm{NH}_{3}\), all bond dipoles add up, creating a large overall dipole moment. In contrast, the bond dipoles in \(\mathrm{NF}_{3}\) partially cancel each other out due to the molecule's geometry, resulting in a smaller overall dipole moment.

Step 1: Identify Structures and Dipole Moments

The \(\mathrm{NH}_{3}\) and \(\mathrm{NF}_{3}\) molecules have trigonal pyramidal structures. In \(\mathrm{NH}_{3}\), hydrogen is electron-poor and nitrogen is electron-rich, creating a dipole moment directed from hydrogen to nitrogen. In \(\mathrm{NF}_{3}\), fluorine is more electronegative than nitrogen, creating individual bond dipole moments directed from nitrogen to fluorine.

Step 2: Explain NH3 Dipole Moment

In \(\mathrm{NH}_{3}\), the individual bond dipole moments all point to the nitrogen atom. Because of the symmetric, pyramidal shape of the molecule, these bond dipoles do not cancel out, resulting in a large overall molecular dipole moment (1.47 debye).

Step 3: Explain NF3 Dipole Moment

In \(\mathrm{NF}_{3}\), although each N-F bond has a strong dipole moment, the geometry of the molecule causes these dipoles to be oriented in such a way that they partially cancel each other out. The resulting overall dipole moment is smaller (0.24 debye).

Key Concepts

Trigonal Pyramidal Structure

The concept of trigonal pyramidal structure is essential when discussing the molecular geometry of compounds like \(\mathrm{NH}_{3}\) and \(\mathrm{NF}_{3}\). This structure is characterized by a central atom bonded to three other atoms, creating a pyramid-like shape with a triangular base. In the case of \(\mathrm{NH}_{3}\) (ammonia), nitrogen is the central atom, bonding with three hydrogen atoms.

A key feature of this geometry is the presence of a lone pair of electrons on the central atom. This lone pair pushes the bonded atoms slightly closer together, giving rise to the pyramid shape.

• The lone pair of electrons results in a bond angle slightly less than the ideal 109.5° seen in a perfect tetrahedral arrangement, usually around 107°.
• The 3-dimensional shape is not flat, allowing for the existence of an overall dipole moment.

The shape of the molecule plays a vital role in determining the orientation of bond dipoles and ultimately affects the overall molecular dipole moment.

Electronegativity

Electronegativity is a measure of an atom's tendency to attract electrons in a chemical bond. It is an important concept when trying to understand why different molecules exhibit different dipole moments.

In ammonia (\(\mathrm{NH}_{3}\)), nitrogen is more electronegative than hydrogen. This means nitrogen pulls the bonding electrons closer to itself, creating a dipole moment from each hydrogen towards nitrogen. Because nitrogen is at the pyramid's apex, all these dipole moments enhance each other, resulting in a significant overall dipole moment.

• In \(\mathrm{NF}_{3}\), the scenario is different. Here, the electronegativity of fluorine exceeds that of nitrogen, leading to dipoles that point from nitrogen towards fluorine.
• The combined effect of these bond polarities, in combination with the molecular geometry, causes a partial cancelation of the dipole moments.

Electronegativity disparities, therefore, influence the direction of bond dipoles, which is crucial for understanding molecular dipole moments.

Bond Dipole Moment

The bond dipole moment is a vector quantity that represents the separation of electrical charge between two atoms in a covalent bond. It points from the less electronegative atom toward the more electronegative atom, indicating the direction in which the electrons are pulled.

In molecules like \(\mathrm{NH}_{3}\) and \(\mathrm{NF}_{3}\), these dipole moments are particularly insightful. For \(\mathrm{NH}_{3}\), individual bond moments point towards the nitrogen, cumulatively reinforcing and producing a larger overall dipole moment of 1.47 debye.

On the other hand, in \(\mathrm{NF}_{3}\), individual bond dipoles are directed from nitrogen to fluorine. Despite having strong bond dipole moments due to the high electronegativity of fluorine, the shape of the molecule leads to a net dipole moment of only 0.24 debye.

• These vector quantities depend not only on electronegativity differences but also significantly on the spatial arrangement of atoms in the molecule.
• In the case of \(\mathrm{NF}_{3}\), although the bond dipoles are strong, they are oriented in such a way within the trigonal pyramidal structure that they partially negate each other.

Understanding bond dipole moments and their interactions helps explain the resultant total dipole moment of a molecule.