`OCIO` bond angle in `CIO_(2)^(-)` is  

Options

A. `111^@`
B. `109.5^@`
C. `106^@`
D. `104.5^@`

Answer

To determine the bond angle in the \( ClO_2^- \) ion, we can follow these steps:

### Step 1: Determine the Lewis Structure
1. **Count the total valence electrons**: Chlorine (Cl) has 7 valence electrons, and each oxygen (O) has 6 valence electrons. Since we have two oxygen atoms and a negative charge, the total is:
\[ 7 + (2 \times 6) + 1 = 20 \text{ valence electrons} \]

2. **Draw the structure**: Place Cl in the center, bonded to two O atoms. One O will be double-bonded, and the other will be single-bonded. The single-bonded O will have a negative charge.

\[ Cl = O \quad O^{-} \]

### Step 2: Identify Lone Pairs
3. **Lone pairs on the central atom**: Chlorine has 2 lone pairs of electrons. The double-bonded oxygen has no lone pairs, while the single-bonded oxygen has three lone pairs.

### Step 3: Determine Hybridization
4. **Hybridization**: The central atom (Cl) has 2 bonding pairs (from the two O atoms) and 2 lone pairs. The hybridization can be determined as follows:
- Total regions of electron density = 2 (bonding pairs) + 2 (lone pairs) = 4
- This corresponds to \( sp^3 \) hybridization.

### Step 4: Calculate the Bond Angle
5. **Ideal bond angle**: For \( sp^3 \) hybridization, the ideal bond angle is 109.5 degrees.

6. **Adjust for lone pairs**: Lone pairs repel more strongly than bonding pairs, which reduces the bond angle. With 2 lone pairs, the bond angle will be less than the ideal angle. The bond angle will decrease due to lone pair-lone pair repulsion, leading to a bond angle of approximately 105 degrees.

### Conclusion
The bond angle in \( ClO_2^- \) is approximately **105 degrees**.

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