Question

Write Lewis structures for the following. Show all resonancestructures where applicable.a. \(\mathrm{NO}_{2}^{-}, \mathrm{NO}_{3}^{-}, \mathrm{N}_{2} \mathrm{O}_{4}\left(\mathrm{N}_{2} \mathrm{O}_{4} \text { exists as } \mathrm{O}_{2} \mathrm{N}-\mathrm{NO}_{2} .\right)\) b. \(\mathrm{OCN}^{-}, \mathrm{SCN}^{-}, \mathrm{N}_{3}^{-}\) (Carbon is the central atom in \(\mathrm{OCN}^{-}\) and \(\mathrm{SCN}^{-} .\) )

Short answer

(a) NO2-: Two resonance structures with N as the central atom, one N-O single bond, one N=O double bond and formal charge -1 on single-bonded O. NO3-: Three resonance structures with N as the central atom, two N=O double bonds, one N-O single bond and formal charge -1 on single-bonded O. N2O4: One structure with two NO2 groups interconnected via O—N—O single bond and N=O double bond. (b) OCN-: One structure with C as the central atom, a C—O single bond with a -1 formal charge on O, and a C=N double bond. SCN-: One structure with C as the central atom, a C—S single bond with a -1 formal charge on S, and a C≡N triple bond. N3-: Two resonance structures, one N—N single bond, one N≡N triple bond, and formal charge -1 on the terminal N with 2 lone pairs.

Step-by-step solution

(a) NO2- Lewis Structure and Resonance

1. Total valence electrons: (5 from N) + (6*2 from O) + 1 extra electron due to negative charge = 18 electrons 2. Central atom: Nitrogen 3. Distribute electrons: One N-O bond (2 electrons), one N=O bond (4 electrons), and 8 electrons as lone pairs (4 on each O) 4. All atoms have fulfilled the octet rule 5. Resonance structures: Two possible structures with a double bond between N and one of the O atoms 6. Formal charges: O atoms with single bonds have -1 charge, no formal charge on N

(a) NO3- Lewis Structure and Resonance

1. Total valence electrons: (5 from N) + (6*3 from O) + 1 extra electron due to negative charge = 24 electrons 2. Central atom: Nitrogen 3. Distribute electrons: One N—O single bond (2 electrons) and two N=O double bonds (8 electrons), and 14 electrons as lone pairs (6 on each O with a single bond, 2 on each O with a double bond) 4. All atoms have fulfilled the octet rule 5. Resonance structures: Three possible structures with a double bond between N and each of the three O atoms 6. Formal charges: O atoms with single bonds have -1 charge, no formal charge on N

(a) N2O4 Lewis Structure and Resonance

1. Total valence electrons: (5*2 from N) + (6*4 from O) = 36 electrons 2. Central atoms: Two Nitrogen atoms 3. Structure: Draw one NO2 molecule with an N—O single bond and an N=O double bond as derived above, attach the N of a second NO2 molecule to the singly bonded O via another single bond. Distribute the remaining 4 electrons on the O atoms. 4. All atoms have fulfilled the octet rule, and formal charges are as discussed for NO2.

(b) OCN- Lewis Structure and Resonance

1. Total valence electrons: (6 from O) + (4 from C) + (5 from N) + 1 extra electron due to negative charge = 16 electrons 2. Central atom: Carbon 3. Distribute electrons: Form a C=N double bond (4 electrons) and a C—O single bond (2 electrons), distribute the remaining 10 electrons as lone pairs (6 on O and 2 on N) 4. All atoms have fulfilled the octet rule 5. No resonance structures 6. Formal charges: -1 on O, no formal charge on C and N

(b) SCN- Lewis Structure and Resonance

1. Total valence electrons: (6 from S) + (4 from C) + (5 from N) + 1 extra electron due to negative charge = 16 electrons 2. Central atom: Carbon 3. Distribute electrons: Form a C≡N triple bond (6 electrons) and a C—S single bond (2 electrons), distribute the remaining 8 electrons as lone pairs (6 on S and 2 on N) 4. All atoms have fulfilled the octet rule 5. No resonance structures 6. Formal charges: -1 on S, no formal charge on C and N

(b) N3- Lewis Structure and Resonance

1. Total valence electrons: (5*3 from N) + 1 extra electron due to negative charge = 16 electrons 2. Arrange atoms linearly: N—N—N 3. Distribute electrons: Form a N≡N triple bond (6 electrons) and a N—N single bond (2 electrons), distribute the remaining 8 electrons as lone pairs (2 on each terminal N and 4 on the central N) 4. All atoms have fulfilled the octet rule except the central N which has an extended octet 5. Resonance structures: Two possible structures with a triple bond between each of the terminal N atoms and the central N atom 6. Formal charges: -1 on the terminal N with 2 lone pairs, no formal charge on the central N and other terminal N

Key concepts

Resonance Structures

Resonance structures are essential in understanding the different possible configurations of electron arrangements for molecules with multiple bonds and atoms. These structures are a way to represent molecules where the electron arrangement can vary, offering a more comprehensive picture of the molecule's configuration.
For instance, take the molecule \( ext{NO}_2^-\). It can be drawn in multiple ways by placing a double bond between nitrogen (N) and one of the oxygen (O) atoms, and representing them as more than one valid structure. These different Lewis structures are the resonance structures of \( ext{NO}_2^-\).

• Each resonance structure consists of the same number of electrons and adheres to the rules of formal charges and the octet rule.
• The true structure of the molecule is a hybrid of these resonance structures, resulting in a balance of different possible electron arrangements.
• Importantly, resonance does not imply that the bonds are constantly switching places; instead, the "real" structure is a blend of the possible arrangements.

Resonance structures help us understand molecules like \( ext{NO}_3^-\) and \( ext{N}_3^-\), where multiple bonding scenarios can help describe the molecule's most stable form.

Formal Charges

Formal charges are crucial for determining the most valid resonance structure. They explain the charge distribution within a molecule, aiding in the prediction of the most stable electron configuration.
Calculating formal charges involves assigning charges to individual atoms based on electron distribution in the Lewis structure:

• Start with the number of valence electrons in an atom when it's in its elemental form.
• Subtract the electrons assigned to the atom in the structure. Each lone pair counts as two electrons, and each bond counts as one electron for the atom.
• The formal charge calculation is: \(# \text{valence electrons} - (# \text{bonds} + # \text{lone pair electrons})\).

In molecules like \( ext{NO}_2^-\) or \( ext{SCN}^-\), formal charges help emphasize the most stable configurations. For example, in \( ext{NO}_2^-\), the formal charge on the nitrogen enables the identification of which resonance structure contributes more to the resonance hybrid.
Understanding and calculating formal charges helps in predicting which resonance structures are more significant, ensuring stability and adherence to other rules like the octet rule.

Octet Rule

The octet rule is a fundamental concept when drawing Lewis structures. It states that atoms tend to combine in ways that they each have eight electrons in their valence shells, resembling the electron configuration of a noble gas, for greater stability.
Understanding how to apply the octet rule is crucial when constructing molecules like \( ext{OCN}^-\) or \( ext{N}_3^-\). Following this rule:

• Begin by arranging the central atom with the rest of the atoms around it.
• Distribute the bonding and lone pair electrons so that most atoms end up with eight electrons around them.
• Certain exceptions exist, like with molecules involving elements from the third period onwards, where they might exceed the octet rule due to accessible d-orbitals.

In specific molecules discussed, albeit the octet rule is closely adhered to, exceptions happen as seen in \( ext{N}_3^-\) where the central nitrogen atom can bear an extended octet, illustrating more flexible electron sharing. Keeping this rule in mind allows for a systematic approach in drawing proper and stable Lewis structures.