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Chapter 9: Problem 90

Which of the following statements concerning \(\mathrm{SO}_{2}\) is (are) true? a. The central sulfur atom is \(s p^{2}\) hybridized. b. One of the sulfur-oxygen bonds is longer than the other(s). c. The bond angles about the central sulfur atom are about \(120^{\circ} .\) d. There are two \(\sigma\) bonds in \(\mathrm{SO}_{2}\) . e. There are no resonance structures for SO.

Short Answer

Expert verified
The true statements concerning SO2 are: a. The central sulfur atom is sp2 hybridized. c. The bond angles about the central sulfur atom are about \(120^{\circ}\). d. There are two \(\sigma\) bonds in SO2.

Step by step solution

01

1. Determine Hybridization

To determine the hybridization of the central sulfur atom, we need to identify the number of electron pairs around it. Sulfur dioxide (SO2) has the following Lewis structure: O=S=O Around the central sulfur atom, there are three electron domains: two double bonds with oxygen atoms and one lone pair of electrons. Since there are three electron domains, the hybridization will be sp2. So, statement a is true.
02

2. Compare Bond Lengths

\ As the molecule is symmetric, and both sulfur-oxygen bonds are double bonds, they will have the same length. Therefore, no sulfur-oxygen bond in SO2 is longer than the other. So, statement b is false.
03

3. Determine Bond Angles

In an sp2 hybridized molecule, the ideal bond angle between the three electron domains is approximately 120 degrees, due to the trigonal planar geometry. So, statement c is true.
04

4. Count Sigma Bonds

When we have double bonds, they always consist of one sigma bond and one pi bond. In the Lewis structure of SO2, we have two double bonds. So, there are two sigma bonds in SO2, one in each S=O bond. So, statement d is true.
05

5. Identify Resonance Structures

Resonance structures exist when a molecule has multiple ways to arrange its electrons and they can be represented by different Lewis structures. For SO2, the electrons are distributed in a way that both double bonds are equivalent, and there's no other way to depict SO2 with different electron arrangements. So, statement e is false. In conclusion, the true statements are a, c, and d.

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Most popular questions from this chapter

The atoms in a single bond can rotate about the internuclear axis without breaking the bond. The atoms in a double and triple bond cannot rotate about the internuclear axis unless the bond is broken. Why?

The oxyanion of nitrogen in which it has the highest oxidation state is the nitrate ion \(\left(\mathrm{NO}_{3}^{-}\right) .\) The corresponding oxyanion of phosphorus is \(\mathrm{PO}_{4}^{3-}\) . The \(\mathrm{NO}_{4}^{3-}\) ion is known but is not very stable. The \(\mathrm{PO}_{3}-\) ion is not known. Account for these differences in terms of the bonding in the four anions.

Sodium can react with oxygen to form sodium peroxide $\left(\mathrm{Na}_{2} \mathrm{O}_{2}\right),\( which is composed of \)\mathrm{Na}^{+}$ and \(\mathrm{O}_{2} 2-\) ions. Potassium can react with oxygen to form potassium superoxide \(\left(\mathrm{KO}_{2}\right)\) which is composed of \(\mathrm{K}^{+}\) and \(\mathrm{O}_{2}-\) ions. Does the peroxide ion or the superoxide ion have the shorter bond length? Explain.

The three NO bonds in \(\mathrm{NO}_{3}^{-}\) are all equivalent in length and strength. How is this explained even though any valid Lewis structure for \(\mathrm{NO}_{3}^{-}\) has one double bond and two single bonds to nitrogen?

Describe the bonding in the \(\mathrm{CO}_{3}^{2-}\) ion using the localized electron model. How would the molecular orbital model describe the \(\pi\) bonding in this species?
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