B. Lewis Structures and Formal Charge I) Three possible Lewis structures for the thiocyanate ion, NCS-, are given below: N≡C=S N=C=S [N≡C-S]- Complete each structure by adding the lone pair electrons. Determine the formal charges of the atoms in each structure. Formal charge can be used to distinguish between competing structures. In general, the following rules apply: i) The sum of all formal charges in a neutral molecule must be zero. ii) The sum of all formal charges in an ion must equal the charge on the ion. iii) Small or zero formal charges on individual atoms are better than larger ones. iv) When formal charge cannot be avoided on an atom, negative charges are better on more electronegative atoms. Decide which Lewis structure is the preferred one and give an explanation below.
The Correct Answer and Explanation is:
To analyze the thiocyanate ion (NCS⁻), we need to determine the formal charges for each atom in the three given Lewis structures:
- N≡C=S
- N=C=S
- [N≡C-S]⁻
Step 1: Drawing the Lewis Structures
Each structure needs lone pairs added to complete the octets of the atoms. Here’s the breakdown for each structure:
- N≡C=S
- Nitrogen (N) has 5 valence electrons, Carbon (C) has 4, and Sulfur (S) has 6.
- In this structure, nitrogen forms a triple bond with carbon and a single bond with sulfur. The lone pairs on sulfur complete its octet, while carbon has a full octet with the triple bond to nitrogen and the single bond to sulfur. Nitrogen has a lone pair.
- N=C=S
- Nitrogen forms a single bond with carbon, and carbon forms a single bond with sulfur. Each atom has lone pairs to complete its octet.
- [N≡C-S]⁻
- In this structure, nitrogen forms a triple bond with carbon, and carbon forms a single bond with sulfur. Since the ion carries a negative charge, we place a lone pair on sulfur to accommodate the extra electron.
Step 2: Formal Charge Calculation
The formal charge (FC) is calculated using the formula: Formal Charge=Valence Electrons−(Lone Pairs+Bonding Electrons2)\text{Formal Charge} = \text{Valence Electrons} – \left(\text{Lone Pairs} + \frac{\text{Bonding Electrons}}{2}\right)Formal Charge=Valence Electrons−(Lone Pairs+2Bonding Electrons)
- N≡C=S
- Nitrogen (N): 5−(2+62)=+15 – (2 + \frac{6}{2}) = +15−(2+26)=+1
- Carbon (C): 4−(0+62)=04 – (0 + \frac{6}{2}) = 04−(0+26)=0
- Sulfur (S): 6−(6+22)=−16 – (6 + \frac{2}{2}) = -16−(6+22)=−1
Formal charges: +1 (N), 0 (C), -1 (S)
- N=C=S
- Nitrogen (N): 5−(2+22)=05 – (2 + \frac{2}{2}) = 05−(2+22)=0
- Carbon (C): 4−(2+42)=04 – (2 + \frac{4}{2}) = 04−(2+24)=0
- Sulfur (S): 6−(4+22)=06 – (4 + \frac{2}{2}) = 06−(4+22)=0
Formal charges: 0 (N), 0 (C), 0 (S)
- [N≡C-S]⁻
- Nitrogen (N): 5−(2+62)=+15 – (2 + \frac{6}{2}) = +15−(2+26)=+1
- Carbon (C): 4−(0+62)=04 – (0 + \frac{6}{2}) = 04−(0+26)=0
- Sulfur (S): 6−(6+22)=−16 – (6 + \frac{2}{2}) = -16−(6+22)=−1
Formal charges: +1 (N), 0 (C), -1 (S)
Step 3: Determining the Preferred Structure
From the formal charge analysis:
- Structure 1 (N≡C=S) has formal charges of +1 on N and -1 on S. This is not ideal because nitrogen is less electronegative than sulfur and should not carry a positive charge.
- Structure 2 (N=C=S) has no formal charges on any atoms. This is the ideal situation and satisfies the formal charge rules.
- Structure 3 ([N≡C-S]⁻) has the negative charge on sulfur, which is more electronegative and thus better suited for the negative charge. However, this structure still leaves nitrogen with a formal charge of +1.
Conclusion
The preferred Lewis structure is Structure 2 (N=C=S) because it has zero formal charges on all atoms. This structure is most stable, as it adheres to the rule of minimizing formal charges.