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Chapter 9: Problem 76

Explain the following: (a) The peroxide ion, \(\mathrm{O}_{2}^{2-}\), has a longer bond length than the superoxide ion, \(\mathrm{O}_{2}^{-}\). (b) The magnetic properties of \(\mathrm{B}_{2}\) are consistent with the \(\pi_{2 \mathrm{p}}\) MOs being lower in energy than the \(\sigma_{2 p}\) MO. (c) The \(\mathrm{O}_{2}^{2+}\) ion has a stronger O- \(O\) bond than \(\mathrm{O}_{2}\) itself.

Short Answer

Expert verified
(a) The peroxide ion (O2^2-) has a longer bond length than the superoxide ion (O2^-) because it has two additional electrons in the antibonding π2p* orbital, which weakens the O-O bond more than the one additional electron in superoxide ion. (b) The magnetic properties of B2 are consistent with the π2p MOs being lower in energy than the σ2p MO because there are two unpaired electrons in the π2p orbitals, giving rise to the magnetic properties. (c) The O2^2+ ion has a stronger O-O bond than O2 because it has two electrons removed from the antibonding π2p orbitals, reducing the antibonding effect and strengthening the O-O bond.

Step by step solution

01

(a) Peroxide ion (O2^2-) and superoxide ion (O2^-) bond lengths

Both ions are formed from the same parent molecule, O2. In O2^2-, two electrons are added, while in O2^-, only one electron is added. In molecular orbital theory, we can represent the electron configurations for these molecules as: - O2: σ1s2 σ1s*2 σ2s2 σ2s*2 σ2p2 π2p4 - O2^2-: σ1s2 σ1s*2 σ2s2 σ2s*2 σ2p2 π2p6 - O2^-: σ1s2 σ1s*2 σ2s2 σ2s*2 σ2p2 π2p5 In peroxide ion, O2^2-, the additional two electrons occupy the antibonding π2p* orbital. This weakens the O-O bond, resulting in a longer bond length compared to O2. In contrast, in superoxide ion, O2^-, only one additional electron occupies the antibonding π2p* orbital, which creates a weaker bond length increase compared to O2^2-.
02

(b) Magnetic properties of B2

The molecular orbital configuration for a B2 molecule is: - B2: σ1s2 σ1s*2 σ2s2 σ2s*2 π2p2 Considering the observed magnetic properties of B2, they are consistent with the π2p MOs being lower in energy than the σ2p MO. This is because two unpaired electrons are present in π2p orbitals, which gives rise to the magnetic properties of B2.
03

(c) Comparison of O2^2+ ion and O2 bond strength

The electron configurations for the O2^2+ ion and O2 molecule are: - O2: σ1s2 σ1s*2 σ2s2 σ2s*2 σ2p2 π2p4 - O2^2+: σ1s2 σ1s*2 σ2s2 σ2s*2 σ2p2 π2p2 When O2 loses two electrons to form O2^2+, the two electrons are removed from the π2p antibonding orbitals. This reduces the antibonding effect, leading to a stronger O-O bond in O2^2+ than in the parent O2 molecule.

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Most popular questions from this chapter

Consider the following \(\mathrm{XF}_{4}\) ions: $\mathrm{PF}_{4}^{-}, \mathrm{BrF}_{4}^{-}, \mathrm{ClF}_{4}^{+},\( and \)\mathrm{AlF}_{4}^{-}$ (a) Which of the ions have more than an octet of electrons around the central atom? (b) For which of the ions will the electron-domain and molecular geometries be the same? (c) Which of the ions will have an octahedral electron-domain geometry? (d) Which of the ions will exhibit a see-saw molecular geometry?

Sulfur tetrafluoride \(\left(\mathrm{SF}_{4}\right)\) reacts slowly with \(\mathrm{O}_{2}\) to form sulfur tetrafluoride monoxide (OSF_4) according to the following unbalanced reaction: $$ \mathrm{SF}_{4}(g)+\mathrm{O}_{2}(g) \longrightarrow \operatorname{OSF}_{4}(g) $$ The \(O\) atom and the four \(\mathrm{F}\) atoms in \(\mathrm{OSF}_{4}\) are bonded to a central S atom. (a) Balance the equation. (b) Write a Lewis structure of \(\mathrm{OSF}_{4}\) in which the formal charges of all atoms are zero. (c) Use average bond enthalpies (Table 8.3) to estimate the enthalpy of the reaction. Is it endothermic or exothermic? (d) Determine the electron-domain geometry of OSF \(_{4}\), and write two possible molecular geometries for the molecule based on this electron-domain geometry. (e) For each of the molecules you drew in part (d), state how many fluorines are equatorial and how many are axial.

(a) Does \(C S_{2}\) have a dipole moment? If so, in which direction does the net dipole point? (b) Does \(\mathrm{SO}_{2}\) have a dipole moment? If so, in which direction does the net dipole point?

(a) Is the molecule \(\mathrm{BF}_{3}\) polar or nonpolar? (b) If you react \(\mathrm{BF}_{3}\) to make the ion \(\mathrm{BF}_{3}^{2-}\), is this ion planar? (c) Does the molecule \(\mathrm{BF}_{2} \mathrm{Cl}\) have a dipole moment?

Indicate the hybridization of the central atom in (a) $\mathrm{H}_{2} \mathrm{~S}$, (d) \(\mathrm{AlI}_{3}\). (b) \(\mathrm{SeF}_{6},(\mathbf{c}) \mathrm{P}(\mathrm{OH})_{3}\)
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