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Chapter 8: Problem 24

Which of the following statements is(are) true? Correct the false statements. a. It is impossible to satisfy the octet rule for all atoms in \(\mathrm{XeF}_{2}\) b. Because \(\mathrm{SF}_{4}\) exists, \(\mathrm{OF}_{4}\) should also exist because oxygen is in the same family as sulfur. c. The bond in \(\mathrm{NO}^{+}\) should be stronger than the bond in \(\mathrm{NO}^{-}\) d. As predicted from the two Lewis structures for ozone, one oxygen-oxygen bond is stronger than the other oxygenoxygen bond.

Short Answer

Expert verified
The corrected statements are: a. The octet rule is satisfied for all atoms in \(\mathrm{XeF}_{2}\). b. The existence of \(\mathrm{SF}_{4}\) does not guarantee the existence of \(\mathrm{OF}_{4}\) as oxygen has a higher electronegativity and typically forms a maximum of two bonds. c. The bond in \(\mathrm{NO}^{-}\) should be stronger than the bond in \(\mathrm{NO}^{+}\) as it contains a triple bond, while \(\mathrm{NO}^{+}\) contains a double bond. d. In ozone, both oxygen-oxygen bonds are equivalent in strength due to resonance, causing them to be a hybrid of single and double bonds.

Step by step solution

01

a. Analyzing the octet rule in \(\mathrm{XeF}_{2}\)

First, we need to draw the Lewis structure of \(\mathrm{XeF}_{2}\). Xenon (Xe) has 8 valence electrons, and Fluorine (F) has 7. Since there are two Fluorine atoms, the total number of valence electrons in the molecule is 8 + 2 * 7 = 22. In the Lewis structure, there will be two single bonds between Xe and two F atoms, which will satisfy the octet rule for the two F atoms. The remaining 18 electrons (3 lone pairs) will be placed around the Xe atom, satisfying the octet rule for the Xe as well. Thus, this statement is false since the octet rule is satisfied for all atoms in \(\mathrm{XeF}_{2}\).
02

b. Comparing \(\mathrm{SF}_{4}\) and \(\mathrm{OF}_{4}\)

Sulfur and oxygen belong to the same family (Group 16); however, this does not guarantee that \(\mathrm{OF}_{4}\) will exist if \(\mathrm{SF}_{4}\) does. Oxygen is more electronegative than sulfur, making it less likely to form a molecule like \(\mathrm{OF}_{4}\). Oxygen typically forms a maximum of two bonds due to its high electronegativity. Therefore, the statement is false.
03

c. Comparing bond strengths in \(\mathrm{NO}^{+}\) and \(\mathrm{NO}^{-}\)

In \(\mathrm{NO}^{+}\), the nitrogen and oxygen atoms share two electrons, forming a double bond, while in \(\mathrm{NO}^{-}\), they share three electrons, resulting in a triple bond. Triple bonds are generally stronger than double bonds because they involve more electron sharing between the atoms. Thus, the statement is false since the bond in \(\mathrm{NO}^{-}\) should be stronger than the bond in \(\mathrm{NO}^{+}\).
04

d. Analyzing the oxygen-oxygen bonds in ozone

Ozone, \(\mathrm{O}_{3}\), has two resonance structures with one double bond and one single bond between the oxygen atoms. However, due to resonance, the actual structure of ozone is a hybrid of these two resonance structures. This means that both oxygen-oxygen bonds are equivalent in terms of bond order and strength. Consequently, the statement is false because the two oxygen-oxygen bonds in ozone are not different in strength since they are both a combination of single and double bonds due to resonance. In summary, all four statements given are false.

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Most popular questions from this chapter

Identify the five compounds of \(\mathrm{H}, \mathrm{N}\), and \(\mathrm{O}\) described as follows. For each compound, write a Lewis structure that is consistent with the information given. a. All the compounds are electrolytes, although not all of them are strong electrolytes. Compounds \(\mathrm{C}\) and \(\mathrm{D}\) are ionic and compound \(\mathrm{B}\) is covalent. b. Nitrogen occurs in its highest possible oxidation state in compounds A and C; nitrogen occurs in its lowest possible oxidation state in compounds \(\mathrm{C}, \mathrm{D}\), and \(\mathrm{E}\). The formal charge on both nitrogens in compound \(\mathrm{C}\) is \(+1\); the formal charge on the only nitrogen in compound \(\mathrm{B}\) is \(0 .\) c. Compounds A and E exist in solution. Both solutions give off gases. Commercially available concentrated solutions of compound A are normally \(16 M\). The commercial, concentrated solution of compound \(\mathrm{E}\) is \(15 \mathrm{M}\). d. Commercial solutions of compound \(\mathrm{E}\) are labeled with a misnomer that implies that a binary, gaseous compound of nitrogen and hydrogen has reacted with water to produce ammonium ions and hydroxide ions. Actually, this reaction occurs to only a slight extent. e. Compound \(\mathrm{D}\) is \(43.7 \% \mathrm{~N}\) and \(50.0 \% \mathrm{O}\) by mass. If compound D were a gas at STP, it would have a density of \(2.86 \mathrm{~g} / \mathrm{L}\). f. A formula unit of compound \(\mathrm{C}\) has one more oxygen than a formula unit of compound D. Compounds \(\mathrm{C}\) and \(\mathrm{A}\) have one ion in common when compound \(\mathrm{A}\) is acting as a strong electrolyte. g. Solutions of compound \(\mathrm{C}\) are weakly acidic; solutions of compound A are strongly acidic; solutions of compounds \(\mathrm{B}\) and \(\mathrm{E}\) are basic. The titration of \(0.726 \mathrm{~g}\) compound \(\mathrm{B}\) requires \(21.98 \mathrm{~mL}\) of \(1.000 M \mathrm{HCl}\) for complete neutralization.

Borazine \(\left(\mathrm{B}_{3} \mathrm{~N}_{3} \mathrm{H}_{6}\right)\) has often been called "inorganic" benzene. Write Lewis structures for borazine. Borazine contains a sixmembered ring of alternating boron and nitrogen atoms with one hydrogen bonded to each boron and nitrogen.

A common trait of simple organic compounds is to have Lewis structures where all atoms have a formal charge of zero. Consider the following incomplete Lewis structure for an organic compound called methyl cyanoacrylate, the main ingredient in Super Glue. Draw a complete Lewis structure for methyl cyanoacrylate in which all atoms have a formal charge of zero.

Use the following data to estimate \(\Delta H_{\mathrm{f}}^{\circ}\) for potassium chloride. $$ \mathrm{K}(s)+\frac{1}{2} \mathrm{Cl}_{2}(g) \longrightarrow \mathrm{KCl}(s) $$ Lattice energy \(-690 . \mathrm{kJ} / \mathrm{mol}\) Ionization energy for \(\mathrm{K}\) \(419 \mathrm{~kJ} / \mathrm{mol}\) Electron affinity of \(\mathrm{Cl} \quad-349 \mathrm{~kJ} / \mathrm{mol}\) Bond energy of \(\mathrm{Cl}_{2}\) \(239 \mathrm{~kJ} / \mathrm{mol}\) \(\begin{array}{ll}\text { Enthalpy of sublimation for } \mathrm{K} & \text { 90. } \mathrm{kJ} / \mathrm{mol}\end{array}\)

Use Coulomb's law, $$ V=\frac{Q_{1} Q_{2}}{4 \pi \epsilon_{0} r}=2.31 \times 10^{-19} \mathrm{~J} \cdot \mathrm{nm}\left(\frac{Q_{1} Q_{2}}{r}\right) $$ to calculate the energy of interaction for the following two arrangements of charges, each having a magnitude equal to the electron charge. a. \(\stackrel{1 \times 10^{-10} \mathrm{~m}}{\longrightarrow(-1) \longleftrightarrow \infty \longrightarrow(+1) \longleftrightarrow 10^{-10} \mathrm{~m}}{\longleftrightarrow} \stackrel{\leftarrow}{\longleftrightarrow}\) b.
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