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Chapter 9: Problem 102

Write Lewis structures for these four isoelectronic species: (a) CO, (b) \(\mathrm{NO}^{+},\) (c) \(\mathrm{CN}^{-},\) (d) \(\mathrm{N}_{2}\). Show formal charges.

Short Answer

Expert verified
The Lewis structures would depict an arrangement where \(CO\) has a single bond with Oxygen carrying a negative formal charge and Carbon a positive one. \(\mathrm{NO}^{+}\) features a single bond with Nitrogen carrying a negative charge and Oxygen a positive one. For \(\mathrm{CN}^{-}\), a single bond is present with Nitrogen showing a negative formal charge. Lastly, \(\mathrm{N}_{2}\) is observed to have a triple bond with no formal charges present on Nitrogen atoms.

Step by step solution

01

Drawing Lewis Structure for \(CO\)

First, calculate the total number of valence electrons available. Carbon has 4 and oxygen has 6. Therefore, there are 10 electrons in total. Draw C and O side by side, and connect them with a straight line (which represents two shared electrons or one covalent bond). Distribute the remaining 8 electrons amongst the atoms in pairs. The oxygen atom will end up with 6 electrons around it, which gives it a negative formal charge, while the carbon atom will hold a positive formal charge.
02

Drawing Lewis Structure for \(\mathrm{NO}^{+}\)

In \(\mathrm{NO}^{+}\), the nitrogen atom provides 5 valence electrons and oxygen provides 6, but since the molecule carries a +1-positive charge, you have to subtract one electron from the total count, leaving us with 10 electrons. After forming a single bond between the N and O atom, which uses up two electrons, distribute the remaining 8 electrons in pairs to each atom. The Nitrogen atom will own five electrons (a negative formal charge) while Oxygen will own a positive charge.
03

Drawing Lewis Structure for \(\mathrm{CN}^{-}\)

In \(\mathrm{CN}^{-}\), carbon contributes 4 valence electrons while nitrogen provides 5. The molecule carries a negative charge, so add an extra electron to the total. There are 10 electrons in total. After single bonding C and N, distribute the remaining 8 electrons giving nitrogen a formal charge of -1.
04

Drawing Lewis Structure for \(\mathrm{N}_{2}\)

For \(\mathrm{N}_{2}\), each Nitrogen atom has 5 valence electrons for a total of 10 electrons. After forming a triple bond between the Nitrogen atoms (using 6 electrons), distribute the remaining 4 electrons in pair to each atom. The Nitrogens atoms remain neutral with no formal charges due to all the electrons being evenly distributed.

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Most popular questions from this chapter

Explain what an ionic bond is.

Write Lewis structures for these molecules: (a) ICl, (b) \(\mathrm{PH}_{3}\) (c) \(\mathrm{P}_{4}\) (each \(\mathrm{P}\) is bonded to three other P atoms), (d) \(\mathrm{H}_{2} \mathrm{~S},\) (e) \(\mathrm{N}_{2} \mathrm{H}_{4}\) (f) \(\mathrm{HClO}_{3},(\mathrm{~g}) \mathrm{COBr}_{2}\) (C is bonded to \(\mathrm{O}\) and \(\mathrm{Br}\) atoms ).

What is lattice energy and what role does it play in the stability of ionic compounds?

Draw Lewis structures of these organic molecules: (a) methanol \(\left(\mathrm{CH}_{3} \mathrm{OH}\right) ;\) (b) ethanol \(\left(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\right)\); (c) tetraethyllead \(\left[\mathrm{Pb}\left(\mathrm{CH}_{2} \mathrm{CH}_{3}\right)_{4}\right]\), which was used in "leaded" gasoline; (d) methylamine \(\left(\mathrm{CH}_{3} \mathrm{NH}_{2}\right)\) (e) mustard gas \(\left(\mathrm{ClCH}_{2} \mathrm{CH}_{2} \mathrm{SCH}_{2} \mathrm{CH}_{2} \mathrm{Cl}\right)\), a poison- ous gas used in World War I; (f) urea \(\left[\left(\mathrm{NH}_{2}\right)_{2} \mathrm{CO}\right]\), a fertilizer; (g) glycine \(\left(\mathrm{NH}_{2} \mathrm{CH}_{2} \mathrm{COOH}\right)\), an amino acid.

Explain how the lattice energy of an ionic compound such as \(\mathrm{KCl}\) can be determined using the Born-Haber cycle. On what law is this procedure based?
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