Question

In \( PO_4^{3-} \), the formal charge on each oxygen atom and the P - O bond order respectively are

• A. -0.75, 0.6
• B. -0.75, 1.0
• C. -0.75, 1.25
• D. -3, 1.25

Hint:

For symmetrical polyatomic ions, formal charge is simply (total charge)/(number of atoms). Bond order is (total valency of central atom)/(number of surrounding atoms) if all surrounding atoms are identical. \

Answer 1

The Correct Option is C

Step 1: Understanding the Concept:
In polyatomic ions where resonance occurs, the charge and the bonds are distributed equally among equivalent atoms. The formal charge per atom and the bond order can be calculated by averaging the total charge and total bonds over the number of surrounding atoms.
Step 2: Key Formula or Approach:
1. Average Formal Charge = \(\frac{\text{Total charge on ion}}{\text{Number of oxygen atoms}}\)
2. Bond Order = \(\frac{\text{Total number of bonds in resonance structures}}{\text{Number of resonating positions}}\)
Step 3: Detailed Explanation:
1. Formal Charge: The total charge on the Phosphate ion (\( PO_{4}^{3-} \)) is \(-3\). There are 4 oxygen atoms. \[ \text{Formal Charge per O} = \frac{-3}{4} = -0.75 \]
2. Bond Order: In the Lewis structure of \( PO_{4}^{3-} \), phosphorus forms one double bond (\( P=O \)) and three single bonds (\( P-O \)) to complete its expanded octet and accommodate the charge. - Total bonds = \( 2 (\text{from double}) + 1 + 1 + 1 = 5 \) bonds. - Number of resonating positions = 4. \[ \text{Bond Order} = \frac{5}{4} = 1.25 \]
Step 4: Final Answer
The formal charge is -0.75 and the bond order is 1.25.