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, the following molecules have net dipole moments: - \(\mathrm{CH}_{2}\mathrm{Cl}_{2}\) (polar) - \(\mathrm{CHCl}_{3}\) (polar) - \(\mathrm{N}_{2}\mathrm{O}\) (polar) - \(\mathrm{NH}_{3}\) (polar) The non-polar molecules are: - \(\mathrm{CCl}_{4}\) (non-polar) - \(\mathrm{CO}_{2}\) (non-polar) - \(\mathrm{PH}_{3}\) (non-polar)

Step-by-step solution

Determine polar bonds

By comparing the electronegativity values of the atoms, we can determine if a molecule has polar bonds or not. C-H bonds are considered non-polar, while C-Cl bonds are polar due to the difference in electronegativity between carbon and chlorine.

Analyze molecular geometry

The structures of the molecules are as follows: - For \(\mathrm{CH}_{2}\mathrm{Cl}_{2}\), the structure is tetrahedral with two hydrogen and two chlorine atoms bonded to the central carbon atom. - For \(\mathrm{CHCl}_{3}\), the structure is also tetrahedral, with one hydrogen and three chlorine atoms bonded to the carbon. - For \(\mathrm{CCl}_{4}\), the structure is a regular tetrahedron, with four chlorine atoms bonded to the carbon.

Calculate the net dipole moment

Based on molecular geometry and bond polarity, we can determine the net dipole moment for each molecule: - \(\mathrm{CH}_{2}\mathrm{Cl}_{2}\): Since it has two polar C-Cl bonds and the molecule has a symmetrical structure, the net dipole moment is non-zero, making the molecule polar. - \(\mathrm{CHCl}_{3}\): The molecule has three polar C-Cl bonds, and due to its geometry, the net dipole moment is also non-zero, making it polar. - \(\mathrm{CCl}_{4}\): Although it has four polar C-Cl bonds, its symmetrical tetrahedral geometry causes the net dipole moment to be zero, making the molecule non-polar. b. \(\mathrm{CO}_{2}, \mathrm{N}_{2}\mathrm{O}\)

Determine polar bonds

C-O and N-O bonds are polar due to the difference in electronegativity between the atoms.

Analyze molecular geometry

The structures of the molecules are as follows: - For \(\mathrm{CO}_{2}\), the structure is linear, with two oxygen atoms double bonded to the central carbon atom. - For \(\mathrm{N}_{2}\mathrm{O}\), the structure is linear, with one nitrogen atom double bonded to another nitrogen atom, which is also single bonded to an oxygen atom.

Calculate the net dipole moment

Based on molecular geometry and bond polarity, we can determine the net dipole moment for each molecule: - \(\mathrm{CO}_{2}\): Although it has two polar C-O bonds, its linear geometry causes the net dipole moment to be zero, making the molecule non-polar. - \(\mathrm{N}_{2}\mathrm{O}\): With one N-O polar bond and a non-polar N-N bond, the molecule has a net dipole moment, making it polar. c. \(\mathrm{PH}_{3}, \mathrm{NH}_{3}\)

Determine polar bonds

P-H bonds are considered non-polar, while N-H bonds are polar due to the difference in electronegativity between nitrogen and hydrogen.

Analyze molecular geometry

The structures of the molecules are as follows: - For \(\mathrm{PH}_{3}\), the structure is trigonal planar, with three hydrogen atoms bonded to the central phosphorus atom. - For \(\mathrm{NH}_{3}\), the structure is a trigonal pyramid, with three hydrogen atoms bonded to the central nitrogen atom and one lone pair on nitrogen.

Calculate the net dipole moment

Based on molecular geometry and bond polarity, we can determine the net dipole moment for each molecule: - \(\mathrm{PH}_{3}\): Since the molecule has non-polar P-H bonds, there is no net dipole moment, making it non-polar. - \(\mathrm{NH}_{3}\): With three polar N-H bonds and an asymmetrical structure, the molecule has a net dipole moment, making it polar.