Three possible Lewis structures for the thiocyanate ion, NCS-, are
(a) Determine the formal charges in each structure.
(b) Based on the formal charges, which Lewis structure is the dominant one?
The Correct Answer and Explanation is :
The thiocyanate ion (NCS-) can be represented by three possible Lewis structures, each involving different bonding arrangements between the nitrogen (N), carbon (C), and sulfur (S) atoms. To determine the formal charges of each structure, we need to use the formal charge formula:
[
\text{Formal charge} = \text{Valence electrons} – \left( \text{Lone electrons} + \frac{1}{2} \text{Bonding electrons} \right)
]
(a) Formal Charges Calculation
Let’s assume the three possible Lewis structures:
- Structure 1: N–C–S
- Nitrogen (N) is double-bonded to carbon, and carbon is single-bonded to sulfur. Sulfur has three lone pairs, nitrogen has one lone pair, and carbon has two bonds.
- Formal charges:
- N: Valence = 5, Lone electrons = 2, Bonding electrons = 4 → Formal charge = 5 – (2 + 4/2) = 0
- C: Valence = 4, Lone electrons = 0, Bonding electrons = 8 → Formal charge = 4 – (0 + 8/2) = 0
- S: Valence = 6, Lone electrons = 6, Bonding electrons = 2 → Formal charge = 6 – (6 + 2/2) = -1
- Structure 2: N≡C–S
- Nitrogen is triple-bonded to carbon, and carbon is single-bonded to sulfur. Nitrogen has one lone pair, carbon has no lone pairs, and sulfur has three lone pairs.
- Formal charges:
- N: Valence = 5, Lone electrons = 2, Bonding electrons = 6 → Formal charge = 5 – (2 + 6/2) = 0
- C: Valence = 4, Lone electrons = 0, Bonding electrons = 6 → Formal charge = 4 – (0 + 6/2) = 0
- S: Valence = 6, Lone electrons = 6, Bonding electrons = 2 → Formal charge = 6 – (6 + 2/2) = -1
- Structure 3: N–C≡S
- Nitrogen is single-bonded to carbon, and carbon is triple-bonded to sulfur. Nitrogen has three lone pairs, carbon has no lone pairs, and sulfur has one lone pair.
- Formal charges:
- N: Valence = 5, Lone electrons = 6, Bonding electrons = 2 → Formal charge = 5 – (6 + 2/2) = +1
- C: Valence = 4, Lone electrons = 0, Bonding electrons = 6 → Formal charge = 4 – (0 + 6/2) = 0
- S: Valence = 6, Lone electrons = 2, Bonding electrons = 6 → Formal charge = 6 – (2 + 6/2) = -1
(b) Dominant Lewis Structure
The dominant Lewis structure is the one with the lowest overall formal charges, as this represents the most stable configuration.
- Structure 1 has a formal charge of 0 on both N and C, and -1 on S, which is relatively stable.
- Structure 2 has the same formal charge distribution as Structure 1, meaning it is also a viable structure.
- Structure 3 places a +1 formal charge on nitrogen, which is less favorable because nitrogen prefers to have a formal charge of 0 or negative, as it is more electronegative than carbon or sulfur.
Thus, Structures 1 and 2 are equally dominant, as they have minimal formal charges. However, structure 2, with a triple bond between nitrogen and carbon, is often considered more favorable due to the increased bond strength.