A student proposes the following Lewis structure for the nitronium NO2+ ion. =N= Assign formal charge to each atom in the student’s Lewis structure. atom formal charge left 0 N right 0
The Correct Answer and Explanation is:
To determine the formal charges of the atoms in the nitronium ion (NO2+), we follow these steps:
Step 1: Determine the structure of NO2+
The nitronium ion (NO2+) consists of a central nitrogen (N) atom bonded to two oxygen (O) atoms. One of the oxygen atoms is double-bonded to nitrogen, while the other is single-bonded to nitrogen and carries a formal positive charge (due to the positive charge on the entire ion). The N-O bond length is shorter in the double bond and longer in the single bond.
Step 2: Calculate the formal charges
The formal charge on an atom is given by the formula:Formal Charge=Valence electrons−(Non-bonding electrons+12×Bonding electrons)\text{Formal Charge} = \text{Valence electrons} – (\text{Non-bonding electrons} + \frac{1}{2} \times \text{Bonding electrons})Formal Charge=Valence electrons−(Non-bonding electrons+21×Bonding electrons)
For the Nitrogen (N) atom:
- Valence electrons: Nitrogen is in Group 15, so it has 5 valence electrons.
- Non-bonding electrons: Nitrogen has no lone pairs in the NO2+ ion.
- Bonding electrons: Nitrogen is bonded to two oxygen atoms. One oxygen has a double bond (4 electrons), and the other has a single bond (2 electrons).
Formal Charge on N=5−(0+12×(4+2))=5−3=+1\text{Formal Charge on N} = 5 – (0 + \frac{1}{2} \times (4 + 2)) = 5 – 3 = +1Formal Charge on N=5−(0+21×(4+2))=5−3=+1
For the oxygen (O) with a double bond:
- Valence electrons: Oxygen is in Group 16, so it has 6 valence electrons.
- Non-bonding electrons: Oxygen has two lone pairs (4 electrons).
- Bonding electrons: Oxygen has a double bond with nitrogen, which contributes 4 electrons.
Formal Charge on O (double-bonded)=6−(4+12×4)=6−6=0\text{Formal Charge on O (double-bonded)} = 6 – (4 + \frac{1}{2} \times 4) = 6 – 6 = 0Formal Charge on O (double-bonded)=6−(4+21×4)=6−6=0
For the oxygen (O) with a single bond:
- Valence electrons: This oxygen also has 6 valence electrons.
- Non-bonding electrons: This oxygen has three lone pairs (6 electrons).
- Bonding electrons: Oxygen has a single bond with nitrogen, contributing 2 electrons.
Formal Charge on O (single-bonded)=6−(6+12×2)=6−7=−1\text{Formal Charge on O (single-bonded)} = 6 – (6 + \frac{1}{2} \times 2) = 6 – 7 = -1Formal Charge on O (single-bonded)=6−(6+21×2)=6−7=−1
Step 3: Assign formal charges
- Nitrogen (N): +1
- Double-bonded Oxygen (O): 0
- Single-bonded Oxygen (O): -1
Step 4: Overall charge of the molecule
The formal charges sum up to give the overall charge of the ion:+1+0+(−1)=+1+1 + 0 + (-1) = +1+1+0+(−1)=+1
Conclusion:
In the correct Lewis structure for the nitronium ion (NO2+), the nitrogen atom carries a formal charge of +1, the double-bonded oxygen carries a formal charge of 0, and the single-bonded oxygen carries a formal charge of -1. These charges are consistent with the overall positive charge on the NO2+ ion.
