Which electron dot structure for OCN\(^-\) places a formal charge of -1 on the most electronegative atom?

• A. O\(:\)C\(riplebond\)N\(:\)^-$ (oxygen with three lone pairs, carbon triple bonded to nitrogen with one lone pair and a negative charge)
• B. O\(:\)C\(-\)N\(:\)^-$ (oxygen with three lone pairs, carbon single bonded to nitrogen with three lone pairs and a negative charge)
• C. O\(:\)C\(=\)N\(:\)^-$ (oxygen with two lone pairs, carbon double bonded to nitrogen with two lone pairs and a negative charge)
• D. O\(:\)^-\(C\)riplebond\(N\):$ (oxygen with three lone pairs and a negative charge, carbon triple bonded to nitrogen with one lone pair)

Answer: (A)

Verified step by step guidance

1. Step 1: Identify the atoms involved and their typical electronegativities. Oxygen (O) is more electronegative than nitrogen (N), and carbon (C) is less electronegative than both.
2. Step 2: Recall that formal charge is calculated using the formula: \(\text{Formal Charge} = \text{Valence Electrons} - \text{Nonbonding Electrons} - \frac{1}{2} \times \text{Bonding Electrons}\).
3. Step 3: For each Lewis structure option, assign electrons to each atom according to the given bonding and lone pairs, then calculate the formal charge on each atom using the formula.
4. Step 4: Determine which structure places the negative formal charge on the most electronegative atom, which is oxygen in this case.
5. Step 5: Compare the formal charges and select the structure where oxygen carries the -1 charge, ensuring the overall charge of the ion is correct and the octet rule is satisfied for all atoms.