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Chapter 10: Problem 78

Antimony pentafluoride, \(\mathrm{SbF}_{5}\), reacts with \(\mathrm{XeF}_{4}\) and \(\mathrm{XeF}_{6}\) to form ionic compounds, \(\mathrm{XeF}_{3}^{+} \mathrm{SbF}_{6}-\) and \(\mathrm{XeF}_{5}^{+} \mathrm{SbF}_{6}^{-} .\) Describe the geometries of the cations and anion in these two compounds.

Short Answer

Expert verified
The geometry of the cations \(XeF_3^+\), \(XeF_5^+\), and the anion \(SbF_6^-\) in the mentioned compounds are trigonal pyramidal, trigonal bipyramidal, and octahedral, respectively.

Step by step solution

01

Identify the geometry of \(XeF_3^+\)

The cation \(XeF_3^+\) has 7 valence electrons on Xe; 6 of these electrons are used in bonding with the three F atoms, while the remaining 1 electron stays as a lone pair. Based on the VSEPR theory, structures that have 4 electron regions (3 bonding pairs and 1 lone pair in this case) adopt a trigonal pyramidal arrangement.
02

Identify the geometry of \(XeF_5^+\)

The cation \(XeF_5^+\) has 7 valence electrons on Xe; all of these electrons are used in bonding with the five F atoms, with no lone pair remaining. According to VSEPR theory, structures that have 5 electron regions (5 bonding pairs and no lone pairs in this case) exhibit a trigonal bipyramidal arrangement.
03

Identify the geometry of \(SbF_6^-\)

The anion \(SbF_6^-\) has 6 valence electrons on Sb, which are all used in bonding with the six F atoms, with no lone pair remaining. Based on VSEPR theory, structures that have 6 electron regions (6 bonding pairs and no lone pairs in this case) adopt an octahedral arrangement.

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Most popular questions from this chapter

The geometries discussed in this chapter all lend themselves to fairly straightforward elucidation of bond angles. The exception is the tetrahedron, because its bond angles are hard to visualize. Consider the \(\mathrm{CCl}_{4}\) molecule, which has a tetrahedral geometry and is nonpolar. By equating the bond moment of a particular \(\mathrm{C}-\mathrm{Cl}\) bond to the resultant bond \(\mathrm{mo}-\) ments of the other three \(\mathrm{C}-\mathrm{Cl}\) bonds in opposite directions, show that the bond angles are all equal to \(109.5^{\circ}\)

Explain why the bond order of \(\mathrm{N}_{2}\) is greater than that of \(\mathrm{N}_{2}^{+}\), but the bond order of \(\mathrm{O}_{2}\) is less than that of \(\mathrm{O}_{2}^{+}\)

\(\mathrm{TCDD},\) or 2,3,7,8 -tetrachlorodibenzo- \(p\) -dioxin, is a highly toxic compound: It gained considerable notoriety in 2004 when it was implicated in the murder plot of a Ukrainian politician. (a) Describe its geometry and state whether the molecule has a dipole moment. (b) How many pi bonds and sigma bonds are there in the molecule?

Use molecular orbital theory to compare the relative stabilities of \(\mathrm{F}_{2}\) and \(\mathrm{F}_{2}^{+}\)

Nitryl fluoride (FNO \(_{2}\) ) is very reactive chemically. The fluorine and oxygen atoms are bonded to the nitrogen atom. (a) Write a Lewis structure for FNO \(_{2}\). (b) Indicate the hybridization of the nitrogen atom. (c) Describe the bonding in terms of molecular orbital theory. Where would you expect delocalized molecular orbitals to form?
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