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Chapter 9: Problem 84

(a) The nitric oxide molecule, NO, readily loses one electron to form the \(\mathrm{NO}^{+}\) ion. Why is this consistent with the electronic structure of \(\mathrm{NO} ?\) (b) Predict the order of the \(\mathrm{N}-\mathrm{O}\) bond strengths in \(\mathrm{NO}, \mathrm{NO}^{+},\) and \(\mathrm{NO}^{-},\) and describe the magnetic properties of each. (c) With what neutral homonuclear diatomic molecules are the \(\mathrm{NO}^{+}\) and \(\mathrm{NO}^{-}\) ions isoelectronic (same number of electrons)?

Short Answer

Expert verified
The electronic structure of NO consists of 11 valence electrons, making it paramagnetic due to one unpaired electron. In contrast, \(\mathrm{NO^{+}}\) has 10 valence electrons, resulting in a completely filled molecular orbitals and diamagnetic behavior. On the other hand, \(\mathrm{NO^{-}}\) has 12 valence electrons and is paramagnetic like NO. The bond strength order for these molecules and ions is NO+ > NO > NO-, due to their respective bond orders of 3, 2.5, and 2. Finally, the ions are isoelectronic with neutral homonuclear diatomic molecules: \(\mathrm{NO^{+}}\) with N2 and \(\mathrm{NO^{-}}\) with O2.

Step by step solution

01

Examine the electronic structure of NO

To determine the electronic structure, we first need to find the number of valence electrons in the molecule. Nitrogen (N) has 5 valence electrons, and oxygen (O) has 6 valence electrons. Thus, the nitric oxide (NO) molecule has a total of 11 valence electrons. To find which orbitals are filled, we can start filling them using the mnemonic (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 5d...): - \(1\sigma_{g} (2e^{⁻})\) - \(1\sigma_{u}^{*} (2e^{⁻})\) - \(2\sigma_{g} (2e^{⁻})\) - \(1\pi_{u} (4e^{⁻})\) - two degenerate orbitals - \(1\pi_{g}^{*} (1e^{⁻})\) - two degenerate orbitals, but only one has an electron
02

Electronic structure of NO+

In the case of NO+, one electron is removed from the molecule, resulting in 10 valence electrons. So, the removal of an electron will be from the highest-energy orbital, which is \(1\pi_{g}^{*}\), thus fully filling all the molecular orbitals.
03

Predict N-O bond strengths in NO, NO+, and NO-

Bond strength is inversely proportional to bond order. For each molecule/ion, we can calculate bond order by considering half the difference of electrons in bonding and anti-bonding orbitals: NO: Bond order = (8 - 3) / 2 = 2.5 NO+: Bond order = (8 - 2) / 2 = 3 NO-: Bond order = (8 - 4) / 2 = 2 Thus, the bond strength follows the order: NO+ > NO > NO-.
04

Describe the magnetic properties of each

A molecule is paramagnetic if it has unpaired electrons and diamagnetic if it has no unpaired electrons. In the case of the electronic structure of NO, there is one unpaired electron present in the highest-energy orbital. Therefore, NO is paramagnetic. For NO+, all molecular orbitals are filled, so there are no unpaired electrons and it is diamagnetic. For NO-, since it has an extra electron, it will also have one unpaired electron and it will be paramagnetic as well.
05

Identify isoelectronic neutral diatomic molecules

To find neutral homonuclear diatomic molecules isoelectronic to the NO+ and NO- ions, we need to find molecules with the same number of valence electrons: NO+: 10 valence electrons - Molecule: N2 (nitrogen) - 2 atoms of nitrogen each with 5 valence electrons -> total of 10 valence electrons NO-: 12 valence electrons - Molecule: O2 (oxygen) - 2 atoms of oxygen each with 6 valence electrons -> total of 12 valence electrons So, the \(\mathrm{NO^{+}}\) ion is isoelectronic with the N2 molecule, and the \(\mathrm{NO^{-}}\) ion is isoelectronic with the O2 molecule.

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Most popular questions from this chapter

Consider the molecule \(\mathrm{PF}_{4}\) Cl. (a) Draw a Lewis structure for the molecule, and predict its electron-domain geometry. (b) Which would you expect to take up more space, a \(\mathrm{P}-\mathrm{F}\) bond or a \(\mathrm{P}-\mathrm{Cl}\) bond? Explain. (c) Predict the molecular geometry of \(\mathrm{PF}_{4} \mathrm{Cl}\). How did your answer for part (b) influence your answer here in part (c)? (d) Would you expect the molecule to distort from its ideal electron-domain geometry? If so, how would it distort?

Consider the Lewis structure for glycine, the simplest amino acid: (a) What are the approximate bond angles about each of the two carbon atoms, and what are the hybridizations of the orbitals on each of them? (b) What are the hybridizations of the orbitals on the two oxygens and the nitrogen atom, and what are the approximate bond angles at the nitrogen? (c) What is the total number of \(\sigma\) bonds in the entire molecule, and what is the total number of \(\pi\) bonds?

(a) What does the term paramagnetism mean? (b) How can one determine experimentally whether a substance is paramagnetic? (c) Which of the following ions would you expect to be paramagnetic: \(\mathrm{O}_{2}^{+}, \mathrm{N}_{2}^{2-}, \mathrm{Li}_{2}^{+}, \mathrm{O}_{2}^{2-} ?\) For those ions that are paramagnetic, determine the number of unpaired electrons.

Sulfur tetrafluoride \(\left(\mathrm{SF}_{4}\right)\) reacts slowly with \(\mathrm{O}_{2}\) to form sulfur tetrafluoride monoxide \(\left(\mathrm{OSF}_{4}\right)\) according to the following unbalanced reaction: $$\mathrm{SF}_{4}(g)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{OSF}_{4}(g)$$ The \(\mathrm{O}\) 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 8.4 ) to estimate the enthalpy of the reaction. Is it endothermic or exothermic? (d) Determine the electron-domain geometry of \(\mathrm{OSF}_{4},\) and write two possible molecular geometries for the molecule based on this electron- domain geometry. (e) Which of the molecular geometries in part (d) is more likely to be observed for the molecule? Explain.

How would you expect the extent of overlap of the bonding atomic orbitals to vary in the series IF, ICl, IBr, and \(I_{2}\) ? Explain your answer.
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