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Chapter 9: Problem 59

Draw three resonance structures for the molecule \(\mathrm{N}_{2} \mathrm{O}_{3}\) (atomic arrangement is \(\mathrm{ONNO}_{2}\) ). Show formal charges.

Short Answer

Expert verified
The completed resonance structures of the molecule ONNO2 would have two additional structures from the base case. The base case has formal charges of 0 on the nitrogen atoms and two oxygen atoms with double bonds, and -1 on the single-bonded oxygen atoms. The resonance structures would display the effects of electron movement.

Step by step solution

01

Start with the basic structure

In this case, the atomic arrangement is ONNO2. Draw this by placing three oxygen atoms around one nitrogen, with nitrogen-nitrogen bonding, and adding another oxygen atom to the remaining nitrogen.
02

Add lone electron pairs and bonds

Every oxygen atom will get six electrons (three lone pairs). The outer Nitrogen gets a lone pair and the inner Nitrogen gets two. Draw double bonds to the two outermost oxygens.
03

Calculate formal charges

The formal charge is calculated as: Formal charge = Valence electrons - (non-bonding electrons + 1/2 bonding electrons). Apply this formula on each atom to get the formal charges.
04

Draw additional resonance structures

For resonance structures, the lone pair of electron on the inner nitrogen can form a double bond with either of the adjacent single-bonded oxygens. Shift the existing double bond to the oxygen, transforming it into a lone pair. Do this first for one oxygen, and then for another, creating two additional resonance structures.
05

Re-calculate formal charges for the additional structures

For each resonance structure drawn in the previous step, re-calculate the formal charges using the same formula used in step 3. This will show how electron distribution affects the formal charge of each atom.

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Most popular questions from this chapter

Comment on the correctness of the statement, "All compounds containing a noble gas atom violate the octet rule."

From the lattice energy of \(\mathrm{KCl}\) in Table 9.1 and the ionization energy of \(\mathrm{K}\) and electron affinity of \(\mathrm{Cl}\) in Tables 8.2 and 8.3 , calculate the \(\Delta H^{\circ}\) for the reaction $$ \mathrm{K}(g)+\mathrm{Cl}(g) \longrightarrow \mathrm{KCl}(s) $$ (a) Draw three resonance structures to represent the ion. (b) Given the following information $$ 2 \mathrm{H}+\mathrm{H}^{+} \longrightarrow \mathrm{H}_{3}^{+} \quad \Delta H^{\circ}=-849 \mathrm{~kJ} / \mathrm{mol} $$ and $$ \mathrm{H}_{2} \longrightarrow 2 \mathrm{H} \quad \Delta H^{\circ}=436.4 \mathrm{~kJ} / \mathrm{mol} $$ calculate \(\Delta H^{\circ}\) for the reaction $$ \mathrm{H}^{+}+\mathrm{H}_{2} \longrightarrow \mathrm{H}_{3}^{+} $$

Hydrazine borane, \(\mathrm{NH}_{2} \mathrm{NH}_{2} \mathrm{BH}_{3}\), has been proposed as a hydrogen storage material. When reacted with lithium hydride (LiH), hydrogen gas is released: $$ \mathrm{NH}_{2} \mathrm{NH}_{2} \mathrm{BH}_{3}+\mathrm{LiH} \longrightarrow \mathrm{LiNH}_{2} \mathrm{NHBH}_{3}+\mathrm{H}_{2} $$ Write Lewis structures for \(\mathrm{NH}_{2} \mathrm{NH}_{2} \mathrm{BH}_{3}\) and \(\mathrm{NH}_{2} \mathrm{NHBH}_{3}^{-}\) and assign all formal charges.

What is Lewis's contribution to our understanding of the covalent bond?

Why does the octet rule not hold for many compounds containing elements in the third period of the periodic table and beyond?
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