Question

Which of the following molecules have net dipole moments? For the molecules that are polar, indicate the polarity of each bond and the direction of the net dipole moment of the molecule. a. \(\mathrm{CH}_{2} \mathrm{Cl}_{2}, \mathrm{CHCl}_{3}, \mathrm{CCl}_{4}\) b. \(\mathrm{CO}_{2}, \mathrm{~N}_{2} \mathrm{O}\) c. \(\mathrm{PH}_{3}, \mathrm{NH}_{3}\)

Short answer

In summary: a. CH$_{2}$Cl$_{2}$ and CHCl$_{3}$ are polar, while CCl$_{4}$ is nonpolar. b. CO$_{2}$ is nonpolar, while N$_{2}$O is polar. c. PH$_{3}$ is nonpolar, while NH$_{3}$ is polar.

Step-by-step solution

Part a) Analyse CH2Cl2, CHCl3, CCl4

Step 1: Determine the shape of each molecule. CH2Cl2: This molecule is tetrahedral in shape. CHCl3: This molecule is also tetrahedral. CCl4: This molecule is tetrahedral as well. Step 2: Determine the polarity of each bond. Electronegativity values: C (2.5), H (2.1), Cl (3.0) Looking at these values, we see that Cl is more electronegative than C, and C is more electronegative than H, so all the bonds in these molecules are polar. Step 3: Determine overall molecular polarity. CH2Cl2: The polar bond dipoles do not cancel each other completely leading to a net dipole moment,thus, it is polar. CHCl3: The bond dipoles do not cancel each other completely, and there is a net dipole moment. So, it is polar. CCl4: Despite the C-Cl bonds being polar, their dipoles cancel out due to symmetry, and thus there is no net dipole moment in this molecule, making it nonpolar.

Part b) Analyse CO2, N2O

Step 1: Determine the shape of each molecule. CO2: This molecule is linear. N2O: This molecule is linear. Step 2: Determine the polarity of each bond. Electronegativity values: C (2.5), O (3.5), N (3.0) There is a difference in electronegativity between atoms, so the bonds here are polar. Step 3: Determine overall molecular polarity. CO2: The two bond dipoles cancel out due to symmetry in the linear structure (O=C=O), so there is no net dipole moment - making the molecule nonpolar despite the polar bonds. N2O: The difference in electronegativity between N and O cause a net dipole moment from the centre N atom to the right, towards the O atom. So, the N2O is polar despite its linear shape.

Part c) Analyse PH3, NH3

Step 1: Determine the shape of each molecule. PH3: This molecule is trigonal pyramidal. NH3: This molecule is also trigonal pyramidal. Step 2: Determine the polarity of each bond. Electronegativity values: P (2.1), H (2.1), N (3.0) The P-H bonds are non-polar, but the N-H bonds are polar because of the electronegativity difference. Step 3: Determine overall molecular polarity. PH3: Although the shape is trigonal pyramidal, there is no difference in electronegativity between P and H, so there's no dipole moment. Thus, the molecule is nonpolar. NH3: The polar bonds and the lack of symmetry (due to the lone pair on nitrogen) mean that the bond dipoles do not cancel out, and there is a net dipole moment from the centre N atom towards the direction of the lone pair. Therefore, the NH3 molecule is polar.

Key concepts

Dipole Moment

The dipole moment of a molecule is a vector quantity that represents the overall polarity of a molecule. It is determined by both the magnitude and direction of individual bond dipoles within the molecule. A net dipole moment occurs when these bond dipoles do not cancel each other out and thus, some part of the molecule has a partial charge. This is important for understanding molecular behavior in various physical and chemical processes.

The dipole moment is influenced by molecular shape and the difference in electronegativity between the atoms involved in bonding. For instance, in molecules like \( ext{CH}_2 ext{Cl}_2\) with a tetrahedral geometry, the dipoles do not cancel out, resulting in a polar molecule. Conversely, in a symmetrical molecule like \( ext{CCl}_4\), the dipoles cancel out completely, making it nonpolar. This highlights how molecular symmetry can affect the overall dipole moment.

Electronegativity

Electronegativity is a measure of an atom's ability to attract and hold onto shared electrons within a bond. It is a key concept in determining the polarity of a bond. The greater the difference in electronegativity between the two bonded atoms, the more polar the bond will be.

For example, in \(\text{CO}_2\), oxygen is more electronegative than carbon. This means that the shared electrons are drawn closer to the oxygen atoms, creating partial charges on both the oxygen and carbon. However, the linear geometry of \(\text{CO}_2\) allows these bond dipoles to cancel out, ultimately making the molecule nonpolar.

In contrast, electronegativity differences in molecules like \(\text{NH}_3\) lead to polar bonds because nitrogen is much more electronegative than hydrogen. This creates a persistent dipole moment and results in an overall polar molecule.

Molecular Geometry

Molecular geometry refers to the spatial arrangement of atoms within a molecule. It is a critical factor in determining whether the individual bond dipoles will cancel out or reinforce each other, affecting the molecule's overall polarity.

Tetrahedral geometries, as seen in \(\text{CH}_2\text{Cl}_2\), can produce a net dipole depending on bond polarity and how the atoms are positioned. The bent angles in nonlinear shapes can prevent dipole cancellation, leading to polar molecules. Conversely, linear geometries can either enhance or nullify dipole moments based on symmetry. For instance, \(\text{CO}_2\) has linear geometry and symmetrical charge distribution, leading to a nonpolar molecule.

The presence of lone pairs, as in \(\text{NH}_3\), can alter geometry from a symmetric ideal, contributing to net polarity of the molecule. Thus, understanding geometry is crucial for predicting molecular behavior and interaction.

Polar Bonds

Polar bonds arise due to differences in electronegativity between two bonded atoms. When one atom has a higher electronegativity, it attracts the shared electrons more strongly, creating a partial negative charge on itself and a partial positive charge on the other atom.

These polar bonds are key to forming molecules with distinctive interactions and behaviors. In \(\text{CHCl}_3\), the disparity in electronegativity between chlorine and hydrogen leads to polar bonds, resulting in a polar molecule overall due to the asymmetric tetrahedral shape. However, it is important to recognize that the presence of polar bonds does not always ensure a polar molecule.

In cases like \(\text{CCl}_4\), despite containing polar \(\text{C—Cl}\) bonds, the molecule remains nonpolar because the tetrahedral symmetry leads to complete cancellation of these dipoles. Therefore, it's essential to consider both individual bond polarity and molecular geometry to accurately determine a molecule's net polarity.