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Chapter 9: Problem 53

Draw three reasonable resonance structures for the \(\mathrm{OCN}^{-}\) ion. Show formal charges.

Short Answer

Expert verified
The resonance structures will have either a double bond between Carbon and Oxygen, a triple bond between Carbon and Nitrogen (with negative charge on Nitrogen), or a double bond between both Carbon and Oxygen, and Carbon and Nitrogen (with negative charge on Carbon).

Step by step solution

01

Draw Lewis structure for the OCN- ion

Start by drawing the Lewis structure for the \(\mathrm{OCN}^{-}\) ion. According to Lewis structure, the molecule is made of Carbon (C), Oxygen (O), Nitrogen (N) and an extra electron due to negative charge. The total number of valence electrons available is 4 (for C) + 6 (for O) + 5 (for N) + 1 (due to -1 charge) = 16 electrons. As Carbon is less electronegative, it is located in the center while Oxygen and Nitrogen on the sides. From those 16 electrons, two will go to the C-N bond and two will go to the C-O bond, and the remaining 12 electrons will be distributed as lone pairs on the O and N atoms.
02

Create resonance structures

A resonance structure can be created by moving electrons around. We have three resonance structures: (i) electrons move from the oxygen to Carbon making double bond (C=O) and moving a pair of electrons from C-N bond to Nitrogen atom (ii) when the structure has a triple bond between carbon and Nitrogen and a single bond between Carbon and Oxygen (iii) the another resonance structure when there's a double bond between Carbon and Nitrogen and a double bond between Carbon and Oxygen. Each of this structure will meet the octet rule for the respective atoms
03

Assign formal charges

In each resonance structure, calculate formal charge for each atom to double check that the overall charge given in the problem (-1) is correct and each atom has full valence electrons. The formal charge is calculated as: Formal charge = valence electrons - (non-bonding electrons + 1/2 bonding electrons). After calculation, you'll find that first structure has a -1 charge on Oxygen, second one has -1 charge on Nitrogen and third one has -1 on Carbon. The overall charge of each resonance structure matches with the overall -1 charge of the \(\mathrm{OCN}^{-}\) ion. None of the resonance is 'more correct' than the others, They simply show the delocalization of the electrons in the ion.

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Most popular questions from this chapter

Draw a Lewis structure for each of these organic molecules in which the carbon atoms are bonded to each other by single bonds: \(\mathrm{C}_{2} \mathrm{H}_{6}, \mathrm{C}_{4} \mathrm{H}_{10}, \mathrm{C}_{5} \mathrm{H}_{12}\).

Classify these bonds as ionic, polar covalent, or \(\mathrm{co}-\) valent, and give your reasons: (a) the CC bond in \(\mathrm{H}_{3} \mathrm{CCH}_{3},\) (b) the KI bond in \(\mathrm{KI},(\mathrm{c})\) the \(\mathrm{NB}\) bond in \(\mathrm{H}_{3} \mathrm{NBCl}_{3},\) (d) the \(\mathrm{ClO}\) bond in \(\mathrm{ClO}_{2}\).

Draw four reasonable resonance structures for the \(\mathrm{PO}_{3} \mathrm{~F}^{2-}\) ion. The central \(\mathrm{P}\) atom is bonded to the three \(\mathrm{O}\) atoms and to the \(\mathrm{F}\) atom. Show formal charges.

(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 \(\mathrm{F}_{2}^{-}\).

Write a Lewis structure for \(\mathrm{SbCl}_{5}\). Is the octet rule obeyed in this molecule?
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