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Chapter 9: Problem 34

For each of the following molecules or ions that contain sulfur, write the Lewis structure(s), predict the molecular structure (including bond angles), and give the expected hybrid orbitals for sulfur. $$ \begin{array}{l}{\text { a. } \mathrm{SO}_{2}} \\ {\text { b. } \mathrm{SO}_{3}}\end{array} $$ $$ \text {c} \mathrm{s}_{2} \mathrm{O}_{3}^{2-}\left[\begin{array}{c}{\mathrm{o}} \\\ {\mathrm{s}-\mathrm{s}-\mathrm{o}} \\ {\mathrm{o}} \\\ {\mathrm{o}}\end{array}\right]^{2-} $$ e. \(\mathrm{SO}_{3}^{2-}\) f. \(\mathrm{SO}_{4}^{2-}\) g. \(\mathrm{SF}_{2}\) h. \(\mathrm{SF}_{4}\) i. \(\mathrm{SF}_{6}\) j. \(\mathrm{F}_{3} \mathrm{S}-\mathrm{SF}\) k. \(\mathrm{SF}_{5}+\)

Short Answer

Expert verified
The molecular structures, bond angles, and hybrid orbitals for each given molecule or ion are as follows: a. \(\mathrm{SO_2}\): Bent, ~119°, sp^2, b. \(\mathrm{SO_3}\): Trigonal planar, 120°, sp^2, c. \(\mathrm{S_2O_3^{2-}}\): Trigonal pyramidal, ~107°, sp^3, e. \(\mathrm{SO_{3}^{2-}}\): Trigonal pyramidal, ~107°, sp^3, f. \(\mathrm{SO_{4}^{2-}}\): Tetrahedral, 109.5°, sp^3, g. \(\mathrm{SF_2}\): Bent, ~98°, sp^3, h. \(\mathrm{SF_4}\): See-Saw, 120° and 90°, sp^3d, i. \(\mathrm{SF_6}\): Octahedral, 90°, sp^3d^2, j. \(\mathrm{F_3S-SF}\): Trigonal pyramidal, ~107°, sp^3, k. \(\mathrm{SF_5^+}\): Trigonal bipyramidal, 120° and 90°, sp^3d.

Step by step solution

01

1. Draw the Lewis Structure for each molecule or ion

To draw the Lewis structure for each molecule or ion, we should find the total number of valence electrons and arrange the atoms such that they achieve a stable state. We can do this by satisfying the octet rule for most of the elements, while sulfur can exceed the octet rule.
02

2. Determine Electron Domain and Molecular Structure

After obtaining Lewis structures, we need to determine the electron domain by counting the number of bonding electron groups and lone pairs around the sulfur atom. From this, we can predict the molecular structure and bond angles.
03

3. Identify the Hybridization of the Central Sulfur Atom

The hybridization of the central Sulfur atom can be determined by analyzing electron domains. The process of hybridization combines atomic orbitals into new hybrid orbitals suitable for sharing electrons. Now, let's go through each molecule/ion: a. \(\mathrm{SO_2}\)
04

Lewis Structure

\[\mathrm{O=S=O}\] with one lone pair on the Sulfur atom.
05

Molecular Structure

Bent, with a bond angle of approximately 119°.
06

Hybrid Orbitals

sp^2 b. \(\mathrm{SO_3}\)
07

Lewis Structure

\[\mathrm{O=S(=O)_2}\] with single bonded O having one lone pair.
08

Molecular Structure

Trigonal planar, with bond angle 120°.
09

Hybrid Orbitals

sp^2 c. \(\mathrm{S_2O_3^{2-}}\)
10

Lewis Structure

$$ \text O \text S - \text S - \text O \text O \\ \text O $$ with one lone pair on each Sulfur atom.
11

Molecular Structure

Central SS bond with each Sulfur having Trigonal pyramidal geometries with bond angles approximately 107°.
12

Hybrid Orbitals

sp^3 for both sulfur atoms. e. \(\mathrm{SO_{3}^{2-}}\)
13

Lewis Structure

\[\mathrm{[O=S(=O)O]^{2-}}\] with one lone pair on the Sulfur atom.
14

Molecular Structure

Trigonal pyramidal, with bond angles approximately 107°.
15

Hybrid Orbitals

sp^3 f. \(\mathrm{SO_{4}^{2-}}\)
16

Lewis Structure

\[\mathrm{[S(=O)_4]^{2-}}\] with no lone pairs on the Sulfur atom.
17

Molecular Structure

Tetrahedral, 109.5° bond angles.
18

Hybrid Orbitals

sp^3 g. \(\mathrm{SF_2}\)
19

Lewis Structure

\[\mathrm{F-S-F}\] with two lone pairs on the Sulfur atom.
20

Molecular Structure

Bent, bond angle approximately 98°.
21

Hybrid Orbitals

sp^3 h. \(\mathrm{SF_4}\)
22

Lewis Structure

\[\mathrm{F-S(=F)_2-F}\] with one lone pair on the Sulfur atom.
23

Molecular Structure

See-Saw shape, 120° and 90° bond angles.
24

Hybrid Orbitals

sp^3d i. \(\mathrm{SF_6}\)
25

Lewis Structure

\[\mathrm{S(=F)_6}\] with no lone pairs on the Sulfur atom.
26

Molecular Structure

Octahedral, 90° bond angles.
27

Hybrid Orbitals

sp^3d^2 j. \(\mathrm{F_3S-SF}\)
28

Lewis Structure

\[\mathrm{F_3S-SF}\] with one lone pair on the Sulfur atoms.
29

Molecular Structure

Both Sulfur atoms have Trigonal pyramidal geometries, bond angles approximately 107°.
30

Hybrid Orbitals

sp^3 for both sulfur atoms. k. \(\mathrm{SF_5^+}\)
31

Lewis Structure

\[\mathrm{F_5S^+}\] with no lone pairs on the Sulfur atom.
32

Molecular Structure

Trigonal bipyramidal, 120° and 90° bond angles.
33

Hybrid Orbitals

sp^3d

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Most popular questions from this chapter

The allene molecule has the following Lewis structure: Must all hydrogen atoms lie the same plane? If not, what is their spatial relationship? Explain.

Show how two 2\(p\) atomic orbitals can combine to form a \(\sigma\) or a \(\pi\) molecular orbital.

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

Consider three molecules: A, \(\mathrm{B},\) and \(\mathrm{C}\) . Molecule \(\mathrm{A}\) has a hybridization of \(s p^{3} .\) Molecule \(\mathrm{B}\) has two more effective pairs (electron pairs around the central atom) than molecule A. Molecule C consists of two \(\sigma\) bonds and two \(\pi\) bonds. Give the molecular structure, hybridization, bond angles, and an example for each molecule.

An unusual category of acids known as superacids, which are defined as any acid stronger than 100\(\%\) sulfuric acid, can be prepared by seemingly simple reactions similar to the one below. In this example, the reaction of anhydrous HF with SbF produces the superacid $\left[\mathrm{H}_{2} \mathrm{F}\right]^{+}\left[\mathrm{SbF}_{6}\right]^{-} :$ $$ 2 \mathrm{HF}(l)+\mathrm{SbF}_{5}(l) \longrightarrow\left[\mathrm{H}_{2} \mathrm{F}\right]^{+}\left[\mathrm{SbF}_{6}\right]^{-}(l) $$ a. What are the molecular structures of all species in this reaction? What are the hybridizations of the central atoms in each species? b. What mass of $\left[\mathrm{H}_{2} \mathrm{F}\right]^{+}\left[\mathrm{SbF}_{6}\right]^{-}$ can be prepared when 2.93 \(\mathrm{mL}\) anhydrous \(\mathrm{HF}\) (density $=0.975 \mathrm{g} / \mathrm{mL} )\( and 10.0 \)\mathrm{mL}\( SbFs (density \)=3.10 \mathrm{g} / \mathrm{mL}$ ) are allowed to react?
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