In the reaction `:Cl_(2)+OH^(-)rarrCl^(-)+ClO_(4)^(-)+H_(2)O` :-

Options

A. Chlorine is oxidised
B. Chlorine is reduced.
C. Chlorine is oxidised as well as reduced.
D. Chlorine is neither oxidised nor reduced.

Explanation

To analyze the redox reaction \( \text{Cl}_2 + \text{OH}^- \rightarrow \text{Cl}^- + \text{ClO}_4^- + \text{H}_2\text{O} \), we will follow these steps:

### Step 1: Identify the oxidation states of chlorine in the reactants and products.
- In \( \text{Cl}_2 \), the oxidation state of chlorine is **0** (since it is in its elemental form).
- In \( \text{Cl}^- \), the oxidation state of chlorine is **-1**.
- In \( \text{ClO}_4^- \), we need to calculate the oxidation state of chlorine:
- Let the oxidation state of Cl be \( x \).
- The equation for the oxidation state is:
\[ x + 4(-2) = -1 \]
\[ x - 8 = -1 \implies x = +7 \]
- Therefore, the oxidation state of chlorine in \( \text{ClO}_4^- \) is **+7**.

### Step 2: Determine the changes in oxidation states.
- Chlorine changes from **0** in \( \text{Cl}_2 \) to **-1** in \( \text{Cl}^- \) (reduction).
- Chlorine changes from **0** in \( \text{Cl}_2 \) to **+7** in \( \text{ClO}_4^- \) (oxidation).

### Step 3: Identify the oxidation and reduction processes.
- **Reduction**: The species that gains electrons (decrease in oxidation state) is \( \text{Cl}_2 \) to \( \text{Cl}^- \).
- **Oxidation**: The species that loses electrons (increase in oxidation state) is \( \text{Cl}_2 \) to \( \text{ClO}_4^- \).

### Step 4: Conclusion
Since chlorine undergoes both oxidation and reduction in this reaction, we conclude that chlorine is both oxidized and reduced.

### Final Answer
The correct option is: **Chlorine is oxidized as well as reduced.**

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