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Chapter 4: Problem 92

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

Expert verified
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

01

1. Determine molecular geometries

For each molecule, determine the molecular geometry. This can be accomplished by considering the arrangement of atoms and lone pairs.
02

2. Electronegativity differences between atoms

Determine the electronegativity difference between the atoms in the molecule. The greater the difference, the more polar the bond.
03

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.
04

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.
05

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.
06

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: \(\mathrm{CH}_{2} \mathrm{Cl}_{2}, \mathrm{CHCl}_{3}, \mathrm{N}_{2}\mathrm{O}, \mathrm{PH}_{3}, \mathrm{NH}_{3}\).

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Most popular questions from this chapter

The three most stable oxides of carbon are carbon monoxide \((\mathrm{CO}),\) carbon dioxide \(\left(\mathrm{CO}_{2}\right),\) and carbon suboxide \(\left(\mathrm{C}_{3} \mathrm{O}_{2}\right) .\) The space-filling models for these three compounds are For each oxide, draw the Lewis structure, predict the molecular structure, and describe the bonding (in terms of the hybrid orbitals for the carbon atoms).

Carbon monoxide (CO) forms bonds to a variety of metals and metal ions. Its ability to bond to iron in hemoglobin is the reason that \(\mathrm{CO}\) is so toxic. The bond carbon monoxide forms to metals is through the carbon atom: $$\mathbf{M}-\mathbf{C} \equiv \mathbf{O}$$ a. On the basis of electronegativities, would you expect the carbon atom or the oxygen atom to form bonds to metals? b. Assign formal charges to the atoms in CO. Which atom would you expect to bond to a metal on this basis? c. In the MO model, bonding MOs place more electron density near the more electronegative atom. (See the HF molecule in Figs. \(4-54\) and \(4-55 .\) ) Antibonding MOs place more electron density near the less electronegative atom in the diatomic molecule. Use the MO model to predict which atom of carbon monoxide should form bonds to metals.

The compound \(\mathrm{NF}_{3}\) is quite stable, but \(\mathrm{NCl}_{3}\) is very unstable (NCl \(_{3}\) was first synthesized in 1811 by P. L. Dulong, who lost three fingers and an eye studying its properties). The compounds \(\mathrm{NBr}_{3}\) and \(\mathrm{NI}_{3}\) are unknown, although the explosive compound \(\mathrm{NI}_{3} \cdot \mathrm{NH}_{3}\) is known. Account for the instability of these halides of nitrogen.

Using the molecular orbital model, write electron configurations for the following diatomic species and calculate the bond orders. Which ones are paramagnetic? Place the species in order of increasing bond length and bond energy. a. \(\mathrm{CN}^{+}\) b. CN c. \(\mathrm{CN}^{-}\)

In the hybrid orbital model, compare and contrast \(\sigma\) bonds with \(\pi\) bonds. What orbitals form the \(\sigma\) bonds and what orbitals form the \(\pi\) bonds? Assume the \(z\) -axis is the internuclear axis.
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