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Chapter 9: Problem 116

In 1999 an unusual cation containing only nitrogen \(\left(\mathrm{N}_{5}^{+}\right)\) was prepared. Draw three resonance structures of the ion, showing formal charges. (Hint: The \(\mathrm{N}\) atoms are joined in a linear fashion.)

Short Answer

Expert verified
The nitrogen-containing cation (\(N_5^+\)) can have three resonance structures. Moving as per the electron flow, in the first structure, the central atom has a positive charge, and the second and fourth Nitrogens have a negative charge. In the second structure, we have the double bond shifting towards the right, which makes the second atom positive and the third atom negative. In the third structure, the double bond again shifts towards the left, causing the fourth atom to become negative and the third atom positive. All three structures confirm to the formal charge principle.

Step by step solution

01

Understanding resonance and formal charges

Resonance structures is used in Chemistry to visualize the structure of a molecule or ion that can't be represented by a single Lewis structure. The formal charge is the charge assigned to an atom in a molecule, assuming that electrons in a chemical bond are equally shared between atoms, regardless of relative electronegativity.
02

Draw the first resonance structure

\(\[\)H-:N::N^+\:_::N^-\:_::N^+\:_::N::H\[\)\] where `:` denotes a lone pair of electrons, `-` denotes a single bond, and `::` denotes a double bond. Calculate the formal charges of each atom: The outer Nitrogen atoms have a formal charge of 0 (5 valence electrons 鈥� 1 unbonded electron 鈥� 1/2 * 4 bonding electrons). The inner Nitrogen atoms have a formal charge of -1 (5 valence electrons 鈥� 5 unbonded electron 鈥� 1/2 * 2 bonding electrons). The Nitrogen atom in the middle has a formal charge of +1, the formal charge of ion is +1, this structure is correct.
03

Draw the second resonance structure

\(\[\)H-:N::N^+\:_::N:::N^-\:_::N-H\[\)\] where `:` denotes a lone pair of electrons, `-` denotes a single bond, and `:::` denotes a triple bond. Calculate the formal charges of each atom: The outer Nitrogen atoms have a formal charge of 0 (5 valence electrons 鈥� 1 unbonded electron 鈥� 1/2 * 4 bonding electrons). The inner Nitrogen atoms have a formal charge of -1 (5 valence electrons 鈥� 5 unbonded electron 鈥� 1/2 * 2 bonding electrons). The Nitrogen atom in the middle has a formal charge of +1, the formal charge of ion is +1, this structure is correct.
04

Draw the third resonance structure

\(\[\)H-:N::N^-\:_::N^+\:_::N::N-H\[\)\] where `:` denotes a lone pair of electrons, `-` denotes a single bond, and `::` denotes a double bond. Calculate the formal charges of each atom: The Nitrogen at either end has a formal charge of 0 (5 valence electrons - 1 unbonded electron 鈥� 1/2 * 4 bonding electrons). The Nitrogen atom next to end Nitrogen atom has a formal charge of -1 (5 valence electrons - 5 unbonded electrons 鈥� 1/2 * 2 bonding electrons). The Nitrogen atom in the middle has a formal charge of +1. The formal charge of ion is +1, this structure is correct.

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Most popular questions from this chapter

(a) From these data: $$ \begin{aligned} \mathrm{F}_{2}(g) \longrightarrow 2 \mathrm{~F}(g) & \Delta H_{\mathrm{rxn}}^{\circ}=156.9 \mathrm{~kJ} / \mathrm{mol} \\ \mathrm{F}^{-}(g) \longrightarrow \mathrm{F}(g)+e^{-} & \Delta H_{\mathrm{rxn}}^{\circ}=333 \mathrm{~kJ} / \mathrm{mol} \\ \mathrm{F}_{2}^{-}(g) \longrightarrow \mathrm{F}_{2}(g)+e^{-} & \Delta H_{\mathrm{rxn}}^{\circ}=290 \mathrm{~kJ} / \mathrm{mol} \end{aligned} $$ calculate the bond enthalpy of the \(\mathrm{F}_{2}^{-}\) ion. (b) Explain the difference between the bond enthalpies of \(\mathrm{F}_{2}\) and \(\mathrm{F}_{2}^{-}\).

Draw Lewis structures of these organic molecules: (a) tetrafluoroethylene \(\left(\mathrm{C}_{2} \mathrm{~F}_{4}\right)\) (b) propane \(\left(\mathrm{C}_{3} \mathrm{H}_{8}\right)\) (c) butadiene \(\left(\mathrm{CH}_{2} \mathrm{CHCHCH}_{2}\right.\) ), (d) propyne \(\left(\mathrm{CH}_{3} \mathrm{CCH}\right),\) (e) benzoic acid \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\right) .\) (Hint: To draw \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\), replace an \(\mathrm{H}\) atom in benzene with a COOH group.)

Define bond length, resonance, and resonance structure.

For each of the following pairs of elements, state whether the binary compound they form is likely to be ionic or covalent. Write the empirical formula and name of the compound: (a) I and \(\mathrm{Cl}\), (b) \(\mathrm{Mg}\) and \(\mathrm{F}\).

Of the noble gases, only \(\mathrm{Kr}, \mathrm{Xe},\) and \(\mathrm{Rn}\) are known to form a few compounds with \(\mathrm{O}\) and/or \(\mathrm{F}\). Write Lewis structures for these molecules: (a) \(\mathrm{XeF}_{2},\) (b) \(\mathrm{XeF}_{4}\) (c) \(\mathrm{XeF}_{6},\) (d) \(\mathrm{XeOF}_{4},\) (e) \(\mathrm{XeO}_{2} \mathrm{~F}_{2}\). In each case Xe is the central atom.
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