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Chapter 10: Problem 8

Suggest reasons why the following do not exist as stable molecules: (a) \(\mathrm{H}_{3} ;\) (b) \(\mathrm{HHe} ;\) (c) \(\mathrm{He}_{2}\); (d) \(\mathrm{H}_{3} \mathrm{O}\).

Short Answer

Expert verified
H3, HHe, He2 and H3O do not exist as stably due to the inability to satisfy the conditions for a full valence electron shell in the molecules. For H3 and H3O there are insufficient electrons to form stable bonds, while for HHe and He2, the He atom already has a full valence shell and thus has no incentive to bond.

Step by step solution

01

Analysis of H3

In H3, each hydrogen atom only has 1 electron in their valence shell. When two hydrogen atoms bond, they achieve full valence shell of 2 electrons. In case of H3, one hydrogen atom would have to share its only electron with two other hydrogen atoms, leaving it insufficient electrons to form a full valence shell. This results in instability of H3, hence it doesn't exist as a stable molecule.
02

Analysis of HHe

In HHe, the helium atom (He) has full valence shell with 2 electrons and doesn't need to bond to become stable. Hydrogen (H) needs 1 more electron to reach a full valence shell, but He doesn't require any more electrons. Therefore, He has no incentive to bond with hydrogen, rendering HHe non-existing as a stable molecule.
03

Analysis of He2

Each helium atom in He2 has full valence shell with 2 electrons and doesn't need to bond to become stable. Therefore, helium atoms (He) have no incentive to bond with each other, rendering He2 nonexistent as a stable molecule.
04

Analysis of H3O

For H3O, oxygen (O) has 6 electrons in its valence shell which can link with two of the hydrogen atoms. However, the third hydrogen atom would have no electrons to form a stable bond with, leading to imbalance and instability. Thus, H3O do not exist as a stable molecule.

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Most popular questions from this chapter

Draw plausible Lewis structures for the following species; use expanded valence shells where necessary. (a) \(\mathrm{Cl}_{2} \mathrm{O} ;\) (b) \(\mathrm{PF}_{3} ;\) (c) \(\mathrm{CO}_{3}^{2-} ;\) (d) \(\mathrm{BrF}_{5}\).

Draw a plausible Lewis structure for the following series of molecules and ions: \((a) \operatorname{SiF}_{6}^{2-} ;\) (b) \(\mathrm{PF}_{5} ;\) (c) \(\mathrm{SF}_{4}\); (d) \(\mathrm{XeF}_{4}\). Describe the electron group geometry and molecular structure of these species.

Draw a plausible Lewis structure for the following series of molecules and ions: (a) \(\mathrm{ClF}_{2}^{-} ;\) (b) \(\mathrm{ClF}_{3}\); (c) \(\mathrm{ClF}_{4}^{-} ;\) (d) ClF \(_{5} .\) Describe the electron group geometry and molecular structure of these species.

Hydrogen azide, \(\mathrm{HN}_{3}\), is a liquid that explodes violently when subjected to physical shock. In the \(\mathrm{HN}_{3}\) molecule, one nitrogen- to-nitrogen bond length is \(113 \mathrm{pm},\) and the other is \(124 \mathrm{pm} .\) The \(\mathrm{H}-\mathrm{N}-\mathrm{N}\) bond angle is \(112^{\circ} .\) Draw Lewis structures and a sketch of the molecule consistent with these facts.

Which of the following molecules would you expect to have a resultant dipole moment \((\mu) ;\) (a) \(\mathrm{F}_{2}\), (b) \(\mathrm{NO}_{2}\) (c) \(\mathrm{BF}_{3},\) (d) \(\mathrm{HBr},\) (e) \(\mathrm{H}_{2} \mathrm{CCl}_{2}\), (f) \(\mathrm{SiF}_{4},\) (g) OCS? Explain.
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