Problem 189 - A Complete Resource Book in Chemistry for JEE Main

Question

Which one of the following pairs of molecules will have permanent dipole moments for both members? (a) \(\mathrm{SiF}_{4}\) and \(\mathrm{NO}_{2}\) (b) \(\mathrm{NO}_{2}\) and \(\mathrm{CO}_{2}\) (c) \(\mathrm{NO}_{2}\) and \(\mathrm{O}_{3}\) (d) \(\mathrm{SiF}_{4}\) and \(\mathrm{CO}_{2}\)

Step-by-Step Solution

Answer

(c) \(\mathrm{NO}_{2}\) and \(\mathrm{O}_{3}\)

Step 1: Determine Molecular Geometry of Each Molecule

Examine the geometry of each molecule to understand if symmetry cancels out dipole moments. For example, \(\mathrm{SiF}_{4}\) has a tetrahedral structure with symmetrical distribution, causing dipole moments to cancel out. \(\mathrm{CO}_{2}\) is linear and nonpolar, while \(\mathrm{NO}_{2}\) and \(\mathrm{O}_{3}\) have bent structures, which results in a net dipole moment.

Step 2: Identify Molecules with Net Dipole Moment

Molecules with a net dipole moment (permanent dipole) are those with an asymmetrical shape that causes an uneven distribution of charge. \(\mathrm{NO}_{2}\) and \(\mathrm{O}_{3}\) both have a net dipole moment because of their bent shapes, causing an uneven distribution of electrons.

Step 3: Match the Given Pairs to Findings

From the list, \(\mathrm{NO}_{2}\) and \(\mathrm{O}_{3}\) are known to have permanent dipole moments due to their geometries. \(\mathrm{SiF}_{4}\) and \(\mathrm{CO}_{2}\) are both symmetrical and do not have permanent dipole moments.

Step 4: Conclusion Based on Analysis

Since both \(\mathrm{NO}_{2}\) and \(\mathrm{O}_{3}\) have permanent dipole moments, the correct answer to which pair of molecules have permanent dipole moments for both members is \(\mathrm{NO}_{2}\) and \(\mathrm{O}_{3}\).

Key Concepts

Dipole Moment

A dipole moment arises in a molecule when there is a separation of positive and negative charges. It is a vector quantity, meaning it has both magnitude and direction, pointing from the positive to the negative charge. In simple terms, it's the measure of how unevenly electrical charge is distributed across a molecule. This unevenness results in a net dipole if a molecule is asymmetrical regarding its charge distribution.

Several factors influence whether a molecule has a dipole moment. These include the electronegativity of the atoms involved and the overall shape of the molecule. When the atoms bonded together have different electronegativities, this often results in polar bonds, contributing to the dipole moment.

• Polar bonds can lead to a permanent dipole moment if the molecule is asymmetrically shaped.
• Symmetrical shapes can cause these polarities to cancel out, resulting in no net dipole moment.

Understanding dipole moments is crucial in predicting the behavior of molecules, such as their interactions, solubility, and even how they might combine to form new compounds.

Asymmetrical Molecules

Asymmetrical molecules are those whose shape is not uniform, causing an uneven distribution of charges across the molecule. This asymmetry is an important factor in whether a molecule will have a net dipole moment.

Molecules with an asymmetrical geometry often have non-zero dipole moments since the lack of symmetry means that the vector components of each bond don't fully cancel out.

• This happens when the molecule has bent or angular shapes, like water ( H_2O) or nitrogen dioxide ( NO).
• The configurations also depend on whether the central atom has lone pairs of electrons, which can further disrupt symmetry.

When assessing molecules for symmetry, consider both the geometry and the bond polarity. Differences in electronegativity between bonded atoms contribute to asymmetry, while lone pairs can add to the polar nature of a molecule.
Being able to determine asymmetry helps in predicting molecular behavior and interactions with other molecules.

Bent Molecular Shape

Bent molecular shape is a type of molecular geometry where the bond angles are less than 180 degrees, causing the molecule to have an angular or V-shape. This particular arrangement often leads to a polarity because it disrupts potential symmetry that might cancel out polarities.

Take, for example, water ( H_2O) and ozone ( O_3). In both examples, the presence of lone pairs on the central atom (oxygen in these cases) forces the bonds to adopt a bent shape.

• These molecules have two unshared electron pairs, which assume more space than bonding pairs, pushing the bonded atoms closer together.
• As a result, the asymmetrical charge distribution results in a molecule with a permanent dipole moment.

Bent shapes are predominantly found in molecules containing a central atom from groups 15 to 17 of the periodic table, typically having lone electron pairs that influence the geometry. Learning about bent molecular shapes provides insight into understanding why certain molecules are polar and how they will behave in different chemical contexts and environments.