Question

On the basis of individual bond polarity and orientation, determine whether each of the following molecules would be polar or nonpolar: a. \(\mathrm{H}_{2} \mathrm{O}\) b. \(\mathrm{I}_{2}\) c. \(\mathrm{CF}_{4}\) d. \(\mathrm{NH}_{3}\) e. \(\mathrm{CO}_{2}\)

Short answer

a. Polar (bent shape) b. Nonpolar (linear) c. Nonpolar (tetrahedral) d. Polar (trigonal pyramidal) e. Nonpolar (linear)

Step-by-step solution

- Determine the Molecular Shape

Identify the molecular geometry of each molecule, which can be found using VSEPR theory.

- Determine Bond Polarity

Look up the electronegativities of the atoms in each molecule to determine the polarity of individual bonds.

- Determine the Overall Molecule Polarity

Consider both the polarity of bonds and the molecular geometry to determine if the molecule is polar or nonpolar.

Key concepts

VSEPR theory

To determine the shape of a molecule, we use Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory states that electron pairs around a central atom will arrange themselves as far apart as possible to minimize repulsion. By assessing both bonding pairs (those used in covalent bonds) and lone pairs (those not involved in bonds), we can predict the geometry of the molecule. For example:

• In \(\text{H}_2\text{O}\), there are two bonding pairs and two lone pairs around the oxygen atom, resulting in a bent shape.

Knowing the shape is crucial for understanding if the individual bond polarities will create an overall molecular polarity.

Bond polarity

Bond polarity arises from differences in electronegativity between two bonded atoms. Electronegativity is a measure of how strongly an atom attracts electrons in a bond. When the difference is significant, one atom will attract electrons more strongly, creating a dipole moment with partial charges: one side partially negative and the other partially positive.
For example:

• In \(\text{H}_2\text{O}\), oxygen is more electronegative than hydrogen, so the \(\text{O-H}\) bonds are polar.
• In contrast, \(\text{I}_2\) has identical atoms with no difference in electronegativity, making the bond nonpolar.

These individual bond polarities must be considered together with the molecular shape to determine the overall polarity of the molecule.

Electronegativity

Electronegativity is a key concept in determining bond polarity. It is the ability of an atom to attract shared electrons in a bond. Higher electronegativity means stronger attraction. The scale commonly used is the Pauling scale. For example:

• Fluorine, the most electronegative element, has a value of about 4.0.
• Hydrogen has a value of approximately 2.1.

This difference leads to polar bonds, as seen in molecules like \(\text{NH}_3\), where nitrogen (3.0) is more electronegative than hydrogen, resulting in polar \(\text{N-H}\) bonds.
Combining electronegativity differences and molecular geometry helps predict if a molecule is polar or nonpolar. For instance, even if \(\text{CF}_4\) has very polar \(\text{C-F}\) bonds, its symmetrical tetrahedral shape renders the molecule nonpolar overall.