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

The molecules with a net dipole moment are: \(\mathrm{CH}_{2}\mathrm{Cl}_{2}\) (C-Cl bond dipoles toward Cl), \(\mathrm{CHCl}_{3}\) (C-Cl bond dipoles toward Cl), \(\mathrm{N}_{2}\mathrm{O}\) (net dipole moment toward the more electronegative O atom), \(\mathrm{PH}_{3}\) (P-H bond dipoles toward P), and \(\mathrm{NH}_{3}\) (N-H bond dipoles toward N). Nonpolar molecules are: \(\mathrm{CCl}_{4}\) and \(\mathrm{CO}_{2}\).

Step-by-step solution

1. Determine molecular geometries

For each molecule, determine the molecular geometry. This can be accomplished by considering the arrangement of atoms and lone pairs.

2. Electronegativity differences between atoms

Determine the electronegativity difference between the atoms in the molecule. The greater the difference, the more polar the bond.

3. Analyze bond polarities against molecular geometry

Using the electronegativity differences and molecular geometries, analyze the bond polarities against the molecular geometry to determine if there is a net dipole moment. If the polar bonds cancel out due to symmetry, there will be no net dipole moment.

A. Analyze \(\mathrm{CH}_{2} \mathrm{Cl}_{2}, \mathrm{CHCl}_{3}, \mathrm{CCl}_{4}\)

- \(\mathrm{CH}_{2}\mathrm{Cl}_{2}\): Tetrahedral molecular geometry. Two polar C-Cl bonds have a larger electronegativity difference than the C-H bonds. The C-Cl bond dipoles won't cancel each other out, resulting in a net dipole moment. - \(\mathrm{CHCl}_{3}\): Tetrahedral molecular geometry. Three polar C-Cl bonds have a larger electronegativity difference than the C-H bond. Asymmetrical arrangement, so net dipole moment. - \(\mathrm{CCl}_{4}\): Tetrahedral molecular geometry. Four polar C-Cl bonds cancel each other out due to symmetry, so there is no net dipole moment.

B. Analyze \(\mathrm{CO}_{2}, \mathrm{N}_{2}\mathrm{O}\)

- \(\mathrm{CO}_{2}\): Linear molecular geometry. Two polar C=O bonds (double bond) cancel each other out due to symmetry, so there is no net dipole moment. - \(\mathrm{N}_{2}\mathrm{O}\): Linear molecular geometry. Electronegativity differences between N-N vs N-O atoms, resulting in a net dipole moment toward the more electronegative O atom.

C. Analyze \(\mathrm{PH}_{3}, \mathrm{NH}_{3}\)

- \(\mathrm{PH}_{3}\): Trigonal pyramidal molecular geometry. The polarity of the P-H bonds does not cancel out due to the unsymmetrical arrangement, resulting in a net dipole moment. - \(\mathrm{NH}_{3}\): Trigonal pyramidal molecular geometry. The polarity of the N-H bonds does not cancel out due to the unsymmetrical arrangement, resulting in a net dipole moment. In summary, the molecules with a net dipole moment are: \(\mathrm{CH}_{2} \mathrm{Cl}_{2}, \mathrm{CHCl}_{3}, \mathrm{N}_{2}\mathrm{O}, \mathrm{PH}_{3}, \mathrm{NH}_{3}\).

Key concepts

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. It plays a crucial role in determining the physical and chemical properties of the molecule, including its polarity and molecular dipole moment. The shapes can vary based on the number of atoms and lone pairs surrounding a central atom. Common geometries include linear, tetrahedral, and trigonal pyramidal.

For instance, in the case of the molecule \(CO_2\), the carbon atom forms a straight line with two oxygen atoms, creating a linear shape. This specific geometry results in the cancelling out of the dipole moments from the two \(C=O\) bonds, leading to a non-polar outcome.

On the other hand, molecules like \(CH_2Cl_2\), have a tetrahedral shape, but the arrangement of different atoms (i.e., H and Cl) causes an asymmetrical distribution. This lack of symmetry leads to a net dipole moment, confirming its polar nature. Understanding molecular geometry can help predict how molecules will interact with others and behave in different conditions.

Electronegativity

Electronegativity is the tendency of an atom to attract electrons towards itself. It is a fundamental concept for understanding the nature of chemical bonds and, subsequently, a molecule's polarity. Atoms with higher electronegativity will strongly attract electrons within a bond while atoms with lower electronegativity will release them easily.

The electronegativity difference between two bonded atoms is key to determining whether a bond is polar, non-polar, or ionic. For example, carbon and hydrogen have similar electronegativity values, leading to mostly non-polar \(C-H\) bonds. Conversely, chlorine's electronegativity is significantly higher than that of carbon, making \(C-Cl\) bonds quite polar.

When forming chemical bonds, these differences have important effects on the properties of molecules. Molecules like \(CHCl_3\) display polar properties due to the disparity in electronegativity between its bonds, influencing not only the dipole moment but also how it interacts with other substances.

Polarity of Bonds

The polarity of chemical bonds arises from differences in electronegativity between bonded atoms. When atoms with different electronegativities form a bond, there is an unequal sharing of electrons. This unequal sharing leads to a separation of charge across the bond making it polar. If the atoms have equal electronegativity, the bond is non-polar due to equal sharing of electrons.

Polarity does not only depend on the bond itself but also on the molecular geometry. For example, the bond between nitrogen and hydrogen \(N-H\) in \(NH_3\) is polar due to nitrogen's higher electronegativity. The trigonal pyramidal shape of \(NH_3\) results in an unsymmetrical distribution of these polar bonds, resulting in a net dipole moment.

To determine the overall polarity of a molecule, one must consider both the polarity of individual bonds and the molecule's geometry. In molecules like \(CCl_4\), despite having polar \(C-Cl\) bonds, the symmetrical tetrahedral geometry cancels out these dipoles, making the molecule non-polar. Understanding the interaction between bond polarity and molecular geometry is essential for predicting a molecule's chemical behavior.