Three resonance structures are possible for the thiocyanate ion, SCN- (a) Draw the three resonance structures. (b) Calculate the formal charge on each atom in each resonance structure. (c) Based on formal charges and electronegativity, predict which resonance structure most closely approximates the bonding in this ion? (d) What are the similarities and differences of bonding in SCN – compared to the bonding in OCN
The Correct Answer and Explanation is:
(a) Resonance Structures for SCN⁻
The thiocyanate ion, SCN⁻, can have three main resonance structures. These structures involve the shifting of the bonding electrons between the sulfur, carbon, and nitrogen atoms. The three resonance structures are as follows:
- Structure 1:
- S=C=N⁻
- A double bond between sulfur and carbon, a single bond between carbon and nitrogen, and a negative charge on nitrogen.
- Structure 2:
- S⁻=C=N
- A negative charge on sulfur, a double bond between sulfur and carbon, and a single bond between carbon and nitrogen.
- Structure 3:
- S=C=N⁻
- A single bond between sulfur and carbon, a double bond between carbon and nitrogen, and a negative charge on carbon.
(b) Formal Charges on Each Atom in the Resonance Structures
The formal charge (FC) on an atom is calculated using the formula:FC=Valence electrons−(Lone electrons+12×Bonding electrons)\text{FC} = \text{Valence electrons} – (\text{Lone electrons} + \frac{1}{2} \times \text{Bonding electrons})FC=Valence electrons−(Lone electrons+21×Bonding electrons)
Structure 1: S=C=N⁻
- Sulfur (S):
- Valence electrons: 6
- Lone electrons: 2
- Bonding electrons: 4 (from the double bond with carbon)
- Formal charge: 6−(2+42)=6−4=+26 – (2 + \frac{4}{2}) = 6 – 4 = +26−(2+24)=6−4=+2
- Carbon (C):
- Valence electrons: 4
- Lone electrons: 0
- Bonding electrons: 6 (2 from the bond with sulfur, 4 from the bond with nitrogen)
- Formal charge: 4−(0+62)=4−3=+14 – (0 + \frac{6}{2}) = 4 – 3 = +14−(0+26)=4−3=+1
- Nitrogen (N):
- Valence electrons: 5
- Lone electrons: 2
- Bonding electrons: 2 (from the bond with carbon)
- Formal charge: 5−(2+22)=5−3=−15 – (2 + \frac{2}{2}) = 5 – 3 = -15−(2+22)=5−3=−1
Structure 2: S⁻=C=N
- Sulfur (S):
- Valence electrons: 6
- Lone electrons: 6
- Bonding electrons: 2 (from the double bond with carbon)
- Formal charge: 6−(6+22)=6−7=−16 – (6 + \frac{2}{2}) = 6 – 7 = -16−(6+22)=6−7=−1
- Carbon (C):
- Valence electrons: 4
- Lone electrons: 0
- Bonding electrons: 4 (2 from the bond with sulfur, 2 from the bond with nitrogen)
- Formal charge: 4−(0+42)=4−2=+24 – (0 + \frac{4}{2}) = 4 – 2 = +24−(0+24)=4−2=+2
- Nitrogen (N):
- Valence electrons: 5
- Lone electrons: 2
- Bonding electrons: 2 (from the bond with carbon)
- Formal charge: 5−(2+22)=5−3=+15 – (2 + \frac{2}{2}) = 5 – 3 = +15−(2+22)=5−3=+1
Structure 3: S=C=N⁻
- Sulfur (S):
- Valence electrons: 6
- Lone electrons: 2
- Bonding electrons: 2 (from the bond with carbon)
- Formal charge: 6−(2+22)=6−3=+16 – (2 + \frac{2}{2}) = 6 – 3 = +16−(2+22)=6−3=+1
- Carbon (C):
- Valence electrons: 4
- Lone electrons: 0
- Bonding electrons: 6 (2 from the bond with sulfur, 4 from the bond with nitrogen)
- Formal charge: 4−(0+62)=4−3=+14 – (0 + \frac{6}{2}) = 4 – 3 = +14−(0+26)=4−3=+1
- Nitrogen (N):
- Valence electrons: 5
- Lone electrons: 2
- Bonding electrons: 2 (from the bond with carbon)
- Formal charge: 5−(2+22)=5−3=−15 – (2 + \frac{2}{2}) = 5 – 3 = -15−(2+22)=5−3=−1
(c) Prediction Based on Formal Charges and Electronegativity
The most favorable resonance structure will be the one where formal charges are minimized, and any negative formal charges are placed on the most electronegative atom. Here’s the analysis of each resonance structure:
- Structure 1 has a +2 formal charge on sulfur, which is unfavorable because sulfur is less electronegative than nitrogen and oxygen. Additionally, the formal charge on carbon is +1, and nitrogen carries a negative charge, which is reasonable because nitrogen is more electronegative than sulfur and carbon. However, the high positive charge on sulfur makes this structure less favorable.
- Structure 2 places the negative charge on sulfur, which is less electronegative than nitrogen, making this structure less favorable despite the +2 charge on carbon.
- Structure 3 places the negative charge on nitrogen, which is the most electronegative atom in the molecule. It minimizes formal charges and gives a more reasonable charge distribution. Therefore, this structure is the most stable and most closely approximates the bonding in SCN⁻.
(d) Comparison of SCN⁻ and OCN⁻ Bonding
Both SCN⁻ and OCN⁻ are isoelectronic (same number of electrons) and exhibit resonance, but their bonding differs due to the different atoms involved:
- In SCN⁻, sulfur is less electronegative than oxygen, so the negative charge in the most stable resonance structure is placed on nitrogen, which is the most electronegative atom. Sulfur often acts as the electron donor.
- In OCN⁻, oxygen is more electronegative than sulfur, and as a result, the negative charge is more likely to reside on oxygen, making it the most electronegative element. This gives OCN⁻ a different charge distribution compared to SCN⁻.
Both ions follow similar bonding principles in terms of resonance and charge delocalization, but the differences in electronegativity lead to distinct resonance structures and charge distributions. Oxygen’s higher electronegativity in OCN⁻ makes it more likely to bear the negative charge compared to sulfur in SCN⁻.
