Question

Determine the net molecular dipole of each species: a) \(\mathrm{BF}_{3}\) b) \(\mathrm{NH}_{3}\)

Short answer

\(\mathrm{BF}_3\) has no net dipole; \(\mathrm{NH}_3\) has a net dipole towards nitrogen.

Step-by-step solution

Analyze the Geometry of the Molecule

Begin by examining the molecular geometry of each compound, which influences the molecular dipole. \(\mathrm{BF}_3\) is trigonal planar, while \(\mathrm{NH}_3\) has a trigonal pyramidal shape.

Consider the Symmetry

For \(\mathrm{BF}_3\), the trigonal planar shape means that the \(\mathrm{B-F}\) bonds are evenly distributed and symmetrically arranged around the central boron atom. This symmetry causes the dipoles to cancel each other out, resulting in no net molecular dipole moment.

Evaluate Bond Dipoles for Asymmetrical Molecules

For \(\mathrm{NH}_3\), the shape is trigonal pyramidal due to the lone pair on the nitrogen atom, creating an asymmetrical structure. The nitrogen atom is more electronegative than the hydrogen atoms, causing a net dipole moment directed towards nitrogen.

Determine Net Molecular Dipole Moments

Based on the symmetrical shape of \(\mathrm{BF}_3\), it has no net molecular dipole moment. \(\mathrm{NH}_3\), with its asymmetrical structure and polar bonds, has a net molecular dipole moment.

Key concepts

Molecular Geometry

Molecular geometry is the three-dimensional arrangement of atoms within a molecule. It is a critical factor in determining the properties and reactivity of the molecule. By understanding the geometry, we can predict molecular behavior such as polarity and the distribution of electronic charges.
For trigonal planar geometry, atoms are arranged in a flat, triangular shape with 120-degree angles between any two bonds. This symmetry affects how bond dipoles, which are vector quantities, combine.
In contrast, a trigonal pyramidal geometry involves a triangular base with a central atom that is raised, like a pyramid. This shape doesn't allow bond dipoles to cancel entirely, affecting the molecule's overall polarity. Thus, the molecular geometry plays a key role in defining whether a molecule has a net dipole moment.

Trigonal Planar

The trigonal planar shape is identified by an atom bonded to three others, lying on the same plane. This geometry results from all the surrounding atoms being evenly spaced around a central atom. Examples like \(\mathrm{BF}_{3}\) demonstrate this structure, with the boron atom at the center and the fluorine atoms spread out symmetrically. Each bond angle in \(\mathrm{BF}_{3}\) is 120 degrees.
Due to its symmetry, the individual bond dipoles cancel out. This happens because each dipole is pointed from boron towards fluorine on all sides of the plane, effectively balancing each other. This results in a nonpolar molecule, with no net molecular dipole moment, despite the polar nature of B-F bonds.

Trigonal Pyramidal

The trigonal pyramidal geometry is characterized by a central atom with three bonded atoms and a lone pair, creating an asymmetrical shape. In the case of \(\mathrm{NH}_{3}\), nitrogen is the central atom with hydrogen atoms at the corners of a pyramid shape and a lone pair completing the geometry.
Due to this lone pair, the shape becomes skewed, preventing any cancellation of bond dipoles. Since nitrogen is more electronegative than hydrogen, bond dipoles point towards the nitrogen, and this creates an overall net dipole moment. The presence of the lone pair accentuates this dipole, resulting in a polar molecule with a measurable dipole moment directed towards the nitrogen atom.

Bond Dipoles

Bond dipoles occur due to the difference in electronegativity between two bonded atoms. Electronegativity reflects an atom's ability to attract shared electrons in a bond. In the molecule, dipoles are vectors that show the direction and magnitude of this charge separation.
For molecules with symmetrical geometries like \(\mathrm{BF}_{3}\), the bond dipoles of B-F bonds, although pronounced individually, cancel out because the polarities point equidistantly in all directions, resulting in a zero net dipole moment.
In molecules like \(\mathrm{NH}_{3}\), the asymmetrical trigonal pyramidal shape leads to bond dipoles that do not cancel entirely. As nitrogen pulls more on the electrons, a net dipole is formed pointing towards it, resulting in an overall polar molecule.