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Chapter 10: Problem 56

Write a plausible Lewis structure for \(\mathrm{NO}_{2}\), and indicate whether the molecule is diamagnetic or paramagnetic. Two \(\mathrm{NO}_{2}\) molecules can join together (dimerize) to form \(\mathrm{N}_{2} \mathrm{O}_{4}\). Write a plausible Lewis structure for \(\mathrm{N}_{2} \mathrm{O}_{4},\) and comment on the magnetic properties of the molecule.

Short Answer

Expert verified
The Lewis structure for \(\mathrm{NO}_{2}\) includes an unpaired electron, rendering the molecule paramagnetic. In the case of \(\mathrm{N}_{2} \mathrm{O}_{4}\), the Lewis structure shows that all electrons are paired, meaning the molecule is diamagnetic.

Step by step solution

01

Lewis structure for \(\mathrm{NO}_{2}\)

Firstly, the total number of electrons in \(\mathrm{NO}_{2}\) is calculated by adding the valence electrons of Nitrogen (5 electrons) and Oxygen (6 electrons) twice, resulting in an odd number, 17 electrons. Therefore, \(\mathrm{NO}_{2}\) molecule will have an unpaired electron. A plausible Lewis structure will have Nitrogen in the center, connected by double bond to one Oxygen, and a single bond to the other Oxygen, with the unpaired electron on the Nitrogen.
02

Magnetic property of \(\mathrm{NO}_{2}\)

Since \(\mathrm{NO}_{2}\) molecule has an unpaired electron as seen from the Lewis structure, it would be paramagnetic.
03

Lewis structure for \(\mathrm{N}_{2} \mathrm{O}_{4}\)

The total number of electrons in \(\mathrm{N}_{2} \mathrm{O}_{4}\) is 32, calculated by adding the valence electrons of 2 Nitrogen (5 electrons each) and 4 Oxygen (6 electrons each). In the Lewis structure of \(\mathrm{N}_{2} \mathrm{O}_{4}\), the two Nitrogens are bonded together and each of them is attached to two Oxygens. Each Nitrogen has a bond with one Oxygen atom and two bonds with the other Oxygen, forming a resonance structure.
04

Magnetic property of \(\mathrm{N}_{2} \mathrm{O}_{4}\)

In the \(\mathrm{N}_{2} \mathrm{O}_{4}\) molecule, there are no unpaired electrons from the Lewis structure, implying that this molecule would be diamagnetic.

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Most popular questions from this chapter

In which of the following species is it necessary to employ an expanded valence shell to represent the Lewis structure: $\mathrm{PO}_{4}^{3-}, \mathrm{PI}_{3}, \mathrm{ICl}_{3}, \mathrm{OSCl}_{2}, \mathrm{SF}_{4}, \mathrm{ClO}_{4} ?$ Explain your choices.

Predict the shapes of the following molecules, and then predict which would have resultant dipolemoments: (a) \(\mathrm{SO}_{2} ;\) (b) \(\mathrm{NH}_{3} ;\) (c) \(\mathrm{H}_{2} \mathrm{S} ;\) (d) \(\mathrm{C}_{2} \mathrm{H}_{4} ;\) (e) \(\mathrm{SF}_{6}\); (f) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\).

The greatest bond length is found in (a) \(\mathrm{O}_{2} ;\) (b) \(\mathrm{N}_{2}\) (c) \(\mathrm{Br}_{2} ;\) (d) BrCl.

Use the \(\mathrm{NH}_{3}\) molecule as an example to explain the difference between molecular geometry and electron-group geometry.

One of the following ions has a trigonal-planar shape: $\mathrm{SO}_{3}^{2-} ; \mathrm{PO}_{4}^{3-} ; \mathrm{PF}_{6}^{-} ; \mathrm{CO}_{3}^{2-} .$ Which ion is it? Explain.
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