In `PO_(4)^(3-)` the formal charge on each O-atom and `P-O` bond order respectively are .

Options

A. `-0.75,0.1`
B. `-0.75, 1.25`
C. `-0.75,0.6`
D. `-3,1.25`

Explanation

To solve the question regarding the formal charge on each oxygen atom and the bond order in the phosphate ion \( \text{PO}_4^{3-} \), we will follow these steps:

### Step 1: Draw the Lewis Structure of \( \text{PO}_4^{3-} \)

1. **Count the total number of valence electrons**:
- Phosphorus (P) has 5 valence electrons.
- Each Oxygen (O) has 6 valence electrons, and there are 4 oxygen atoms: \( 4 \times 6 = 24 \).
- The ion has a charge of -3, which means we add 3 more electrons.
- Total = \( 5 + 24 + 3 = 32 \) valence electrons.

2. **Arrange the atoms**:
- Place phosphorus in the center and surround it with the four oxygen atoms.

3. **Distribute the electrons**:
- Start by forming single bonds between phosphorus and each oxygen atom. This uses 8 electrons (4 bonds).
- Distribute the remaining electrons to satisfy the octet rule for oxygen atoms. Each oxygen needs 8 electrons, so we add lone pairs to the oxygen atoms.

4. **Form double bonds if necessary**:
- To minimize formal charges, we can form one double bond between phosphorus and one of the oxygen atoms. This will lead to a structure with one double bond and three single bonds.

### Step 2: Calculate the Formal Charge on Each Oxygen Atom

The formula for formal charge (FC) is:
\[ \text{FC} = \text{Valence electrons} - \left(\text{Non-bonding electrons} + \frac{1}{2} \times \text{Bonding electrons}\right) \]

- For the double-bonded oxygen:
- Valence electrons = 6
- Non-bonding electrons = 4 (2 lone pairs)
- Bonding electrons = 4 (2 bonds)
- FC = \( 6 - (4 + \frac{1}{2} \times 4) = 6 - (4 + 2) = 0 \)

- For each single-bonded oxygen:
- Valence electrons = 6
- Non-bonding electrons = 6 (3 lone pairs)
- Bonding electrons = 2 (1 bond)
- FC = \( 6 - (6 + \frac{1}{2} \times 2) = 6 - (6 + 1) = -1 \)

Thus, the formal charge on the double-bonded oxygen is 0, and on each of the three single-bonded oxygens, it is -1.

### Step 3: Calculate the Bond Order Between Phosphorus and Oxygen

Bond order is calculated using the formula:
\[ \text{Bond Order} = \frac{\text{Total number of bonds}}{\text{Number of resonating structures}} \]

- Total number of bonds in the resonance structures:
- There is 1 double bond and 3 single bonds, giving a total of 5 bonds.

- The number of significant resonance structures for \( \text{PO}_4^{3-} \) is 4 (one for each oxygen atom that can have a double bond with phosphorus).

Thus, the bond order is:
\[ \text{Bond Order} = \frac{5}{4} = 1.25 \]

### Final Results

- **Formal charge on each oxygen atom**:
- One oxygen has a formal charge of 0 (double-bonded).
- Three oxygens have a formal charge of -1 each.

- **Bond order between phosphorus and oxygen**:
- The bond order is 1.25.