Question

Identify each of the following molecules as polar or nonpolar: a. HBr b. \(\mathrm{NF}_{3}\) c. \(\mathrm{CHF}_{3}\)

Short answer

a. Polar, b. Polar, c. Polar.

Step-by-step solution

Determine Molecular Geometry

For each molecule, first draw its Lewis structure to determine the molecular geometry. For example, HBr is linear, \(\text{NF}_3\) is trigonal pyramidal, and \(\text{CHF}_3\) is tetrahedral.

Evaluate Bond Polarities

Check the electronegativity differences between the atoms in each molecule. A significant difference indicates polar bonds. HBr has a polar bond due to the difference between H and Br. In \(\text{NF}_3\), N-F bonds are polar. In \(\text{CHF}_3\), all C-H and C-F bonds are polar.

Assess Molecular Dipoles

For each molecule, determine if the bond polarities cancel each other out or if they result in a dipole moment. HBr, being linear with one polar bond, is polar. \(\text{NF}_3\) has a net dipole moment because of its trigonal pyramidal shape, making it polar. \(\text{CHF}_3\) has polar bonds and does not have a symmetrical shape to cancel out dipoles, making it polar.

Key concepts

molecular geometry

Understanding molecular geometry is key to determining the shape of a molecule. The shape affects how the molecules interact and its overall properties.
The geometry is determined by the number of electron pairs around the central atom. These electron pairs can be bonding pairs or lone pairs.
Here are common geometries:

• Linear: Two regions of electron density, like in HBr, align in a straight line (180° bond angle).
• Trigonal Pyramidal: In \(\text{NF}_{3}\), three bonds and one lone pair create a pyramid shape (107° bond angle).
• Tetrahedral: Found in \(\text{CHF}_{3}\), four regions spread out from the center (109.5° bond angle).

To find the molecular geometry, first draw the Lewis structure. Then, consider both the bonding and lone pairs. This will give you a clear visualization of the molecule's shape.

electronegativity

Electronegativity is a measure of how strongly an atom attracts electrons in a bond. Different atoms have different electronegativity values.

• Higher electronegativity means a greater ability to attract electrons.
• Fluorine has the highest electronegativity of any element.
• If the difference in electronegativity between two atoms is large, the bond is likely to be polar.

For example, in HBr, Br is more electronegative than H, creating a polar bond. Similarly, in \(\text{NF}_{3}\), the N-F bonds are polar due to the difference in electronegativity between nitrogen and fluorine.
In the case of \(\text{CHF}_{3}\), both carbon-hydrogen and carbon-fluorine bonds are polar. Electronegativity values help predict whether a bond will be polar and how electrons are shared.

dipole moment

A dipole moment arises when there is a separation of electric charges. It occurs in polar molecules due to unequal sharing of electrons.

• If a molecule has a net dipole moment, it's polar.
• The shape of the molecule determines whether the dipoles cancel out or not.

In HBr, the polar H-Br bond means the molecule has a dipole moment and is polar.
For \(\text{NF}_{3}\), the trigonal pyramidal shape prevents the dipoles from canceling out, resulting in a net dipole moment and a polar molecule.
Even though \(\text{CHF}_{3}\) is tetrahedral, the electronegativity differences lead to a net dipole moment, making it polar as well.
Understanding dipole moments helps in predicting molecular behavior and intermolecular interactions.