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

Answer

To analyze the given reaction and determine the oxidation states of chlorine, we can follow these steps:

### Step 1: Write down the reaction
The reaction provided is:
\[ \text{Cl}_2 + \text{OH}^- \rightarrow \text{Cl}^- + \text{ClO}_4^- + \text{H}_2\text{O} \]

### Step 2: Identify the oxidation states of chlorine in reactants and products
1. **In Cl2**: The oxidation state of chlorine (Cl) in Cl2 is 0 because it is in its elemental form.
2. **In OH⁻**: The hydroxide ion does not contain chlorine, so we do not need to consider its oxidation state.
3. **In Cl⁻**: The oxidation state of chlorine in Cl⁻ is -1.
4. **In ClO₄⁻**: To find the oxidation state of chlorine in ClO₄⁻, we can use the formula:
\[ \text{Oxidation state of Cl} + 4 \times (\text{Oxidation state of O}) = \text{Charge of the ion} \]
The oxidation state of oxygen (O) is -2. Therefore:
\[ x + 4(-2) = -1 \implies x - 8 = -1 \implies x = +7 \]
So, the oxidation state of chlorine in ClO₄⁻ is +7.

### Step 3: Determine the changes in oxidation states
- Chlorine changes from 0 in Cl₂ to -1 in Cl⁻ (reduction).
- Chlorine changes from 0 in Cl₂ to +7 in ClO₄⁻ (oxidation).

### Step 4: Conclusion on oxidation and reduction
- Since chlorine is being oxidized (from 0 to +7) and reduced (from 0 to -1), we conclude that both oxidation and reduction are occurring in this reaction.

### Final Answer
Chlorine is both oxidized and reduced in this reaction.

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