Two resonance structures for the no2 nitrite ion using an ono arrangement are discussed in the introduction draw two
resonance structures for the nitrite ion using an oon arrangement show all of your workbased on the formal charges on each
atom

Question

Two resonance structures for the NO2- (nitrite ion) using an O-N-O arrangement are discussed in the introduction. Draw two resonance structures for the nitrite ion using an O-O-N arrangement. Show all of your work.Based on the formal charges on each atom, which set of resonance structures seem more reasonable, those with an O-N-O arrangement or an O-O-N arrangement? Explain.Which set of resonance structures consistently provide a negative formal charge on the most electronegative atom?

Andrea · Answer

Let’s break down this question step by step to both draw and analyze the resonance structures for the nitrite ion (NO₂⁻) using both the normal O–N–O arrangement and the hypothetical O–O–N arrangement. We will focus on formal charges and electronegativity to determine which is more reasonable.

✅ Step 1: O–N–O Resonance Structures (Actual Structure)

Structure 1:

   ..
O=N–O⁻
   ..

One oxygen is double bonded to nitrogen.
The other oxygen is single bonded to nitrogen and carries a negative charge.
Nitrogen has a formal charge of 0.
The negative charge resides on the oxygen, which is more electronegative than nitrogen.

Structure 2:

   ..
⁻O–N=O
     ..

This is simply the resonance of the first one, with the positions of the double bond and negative charge swapped.
Again, the negative charge is on oxygen, which is good.

➡️ Both structures have a negative charge localized on oxygen, which is more electronegative, and no unreasonable charges.

✅ Step 2: Hypothetical O–O–N Resonance Structures

This is not the standard bonding order for nitrite, but for the purpose of comparison, let’s draw two hypothetical resonance structures:

Structure 1:

O–O=N

Try assigning formal charges:

Left O: single bond to O, lone pairs → Formal charge = -1
Middle O: single bond to O, double bond to N → Formal charge = 0
N: double bond to O only → Formal charge = 0

But this doesn’t account for all 18 valence electrons of NO₂⁻, and more importantly, this structure has an O–O bond, which is less stable and rare in such ions.

Alternatively:

Structure 2:

⁻O=O–N

This would imply a peroxide-like structure (O=O) bonded to nitrogen, again requiring formal charges and octet expansions or violations, making it less reasonable.

➡️ These O–O–N structures tend to place the negative charge on oxygen, which is okay, but they also imply unstable O–O bonds, which are rare and high energy, especially in resonance-stabilized systems like nitrite.

✅ Step 3: Formal Charge Comparison

Let’s evaluate based on formal charges:

Structure Type	Negative Charge on O?	Stable Bonding?	Formal Charge on N?	O–O Bond Present?	Reasonable?
O–N–O	✅ Yes (most of the time)	✅ Yes (common pattern)	✅ 0	❌ No	✅ Yes
O–O–N	✅ Yes (sometimes)	❌ No (O–O bonds)	✅ or ❌	✅ Yes	❌ No

✅ Final Answers

1. Which set of resonance structures is more reasonable?

The O–N–O set is more reasonable. It avoids unstable O–O bonds and distributes the negative charge on the more electronegative oxygen atoms.

2. Which set consistently places the negative charge on the most electronegative atom?

Both sets place the negative charge on oxygen. However, the O–N–O arrangement does so in a more stable way (without requiring unusual bonds).

The O–N–O resonance structures are more reasonable because:
- They avoid O–O bonds.
- They maintain complete octets.
- They place the negative charge on oxygen.
The O–O–N structures are hypothetical, involve unstable O–O bonding, and are less favored energetically.

Would you like me to draw these structures digitally or as Lewis structures in image format?

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