In `PO_(4)^(3-)` ion the formal charge on the oxygen atom of P-O bond is

Options

A. `+1`
B. `-1`
C. `- 0.75`
D. `+0.75`

Explanation

To determine the formal charge on the oxygen atom in the phosphate ion \( \text{PO}_4^{3-} \), we will follow these steps:

### Step 1: Identify the Valence Electrons
First, we need to know the number of valence electrons for the oxygen atom. Oxygen has 6 valence electrons.

### Step 2: Draw the Lewis Structure of \( \text{PO}_4^{3-} \)
In the phosphate ion, phosphorus (P) is the central atom bonded to four oxygen (O) atoms. One of the oxygen atoms is doubly bonded to phosphorus, while the other three are singly bonded and carry a negative charge. Each singly bonded oxygen atom has three lone pairs of electrons.

### Step 3: Count Non-bonding Electrons
For each oxygen atom in the phosphate ion:
- The oxygen atom that is doubly bonded to phosphorus has 2 lone pairs (4 non-bonding electrons).
- Each of the three singly bonded oxygen atoms has 3 lone pairs (6 non-bonding electrons) and carries a negative charge.

### Step 4: Count Bonding Electrons
In the phosphate ion:
- The doubly bonded oxygen contributes 2 electrons to bonding.
- Each of the three singly bonded oxygen atoms contributes 1 bond (2 electrons) to bonding.

### Step 5: Apply the Formal Charge Formula
The formula for calculating formal charge (FC) is:
\[ \text{FC} = V - N - \frac{B}{2} \]
Where:
- \( V \) = number of valence electrons
- \( N \) = number of non-bonding electrons
- \( B \) = number of bonding electrons

### Step 6: Calculate Formal Charge for Each Oxygen Atom
For the singly bonded oxygen atoms:
- \( V = 6 \) (valence electrons)
- \( N = 6 \) (non-bonding electrons)
- \( B = 2 \) (bonding electrons)

Substituting into the formula:
\[ \text{FC} = 6 - 6 - \frac{2}{2} = 6 - 6 - 1 = -1 \]

Thus, the formal charge on each of the three singly bonded oxygen atoms is -1.

### Step 7: Conclusion
The formal charge on the oxygen atom of the P-O bond in the phosphate ion \( \text{PO}_4^{3-} \) is -1.

---