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Chapter 9: Problem 84

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 degrees. d. There are two \(\sigma\) bonds in \(\mathrm{SO}_{2}\). e. There are no resonance structures for \(\mathrm{SO}_{2}\).

Short Answer

Expert verified
The true statements about \(\mathrm{SO}_{2}\) are: a. The central sulfur atom is \(s p^{2}\) hybridized, c. The bond angles about the central sulfur atom are about 120 degrees, and d. There are two \(\sigma\) bonds in \(\mathrm{SO}_{2}\).

Step by step solution

01

Drawing the Lewis structure and determining the hybridization

: To analyze the molecular structure of \(\mathrm{SO}_{2}\), we need to first draw the Lewis structure. O=S=O The sulfur atom has a double bond with each oxygen atom, and there is a lone pair of electrons on it. The sulfur atom is surrounded by three electron groups (two double bonds and one lone pair). According to the VSEPR theory, with three electron groups, the central atom exhibits \(\mathrm{sp^2}\) hybridization.
02

Checking bond length

: In \(\mathrm{SO}_{2}\), the sulfur atom has double bonds with the two oxygen atoms, and the bond lengths should be the same. Therefore, one bond is not longer than the other.
03

Finding bond angles

: With three electron groups around the central atom sulfur in \(\mathrm{SO}_{2}\), the ideal bond angle should be 120 degrees.
04

Identifying sigma bonds

: In a double bond, there is one sigma bond and one pi bond. Since \(\mathrm{SO}_{2}\) has two double bonds, we can conclude that it has two sigma bonds.
05

Checking for resonance structures

: There are no valid resonance structures for \(\mathrm{SO}_{2}\) where we can draw a different arrangement of the double bonds or preserve the octet rule. Based on our analysis, the true statements are: a. The central sulfur atom is \(s p^{2}\) hybridized. c. The bond angles about the central sulfur atom are about 120 degrees. d. There are two \(\sigma\) bonds in \(\mathrm{SO}_{2}\).

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Most popular questions from this chapter

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. CO b. \(\mathrm{CO}^{+}\) c. \(\mathrm{CO}^{2+}\)

Why must all six atoms in \(\mathrm{C}_{2} \mathrm{H}_{4}\) lie in the same plane?

Values of measured bond energies may vary greatly depending on the molecule studied. Consider the following reactions: $$ \begin{array}{cc} \mathrm{NCl}_{3}(g) \longrightarrow \mathrm{NCl}_{2}(g)+\mathrm{Cl}(g) & \Delta H=375 \mathrm{~kJ} / \mathrm{mol} \\ \mathrm{ONCl}(g) \longrightarrow \mathrm{NO}(g)+\mathrm{Cl}(g) & \Delta H=158 \mathrm{~kJ} / \mathrm{mol} \end{array} $$ Rationalize the difference in the values of \(\Delta H\) for these reactions, even though each reaction appears to involve only the breaking of one \(\mathrm{N}-\mathrm{Cl}\) bond. (Hint: Consider the bond order of the NO bond in ONCl and in NO.)

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?

Which is the more correct statement: "The methane molecule \(\left(\mathrm{CH}_{4}\right)\) is a tetrahedral molecule because it is \(s p^{3}\) hybridized" of "The methane molecule \(\left(\mathrm{CH}_{4}\right)\) is \(s p^{3}\) hybridized because it is a tetrahedral molecule"? What, if anything, is the difference between these two statements?
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