Which of the following is`//` are disproportionation reactions?

Options

A. `2O_(3)rarr3O_(2)`
B. `4KClO_(3)rarr3KClO_(4)+KCl`
C. `2H_(2)O_(2)rarr2H_(2)O+O_(2)`
D. `2KO_(2)+2H_(2)Orarr4KOH+3O_(2)`

Explanation

To determine which of the given reactions are disproportionation reactions, we need to identify if any element in the reaction undergoes both oxidation and reduction simultaneously. Let's analyze each reaction step by step.

### Step 1: Analyze the First Reaction
**Reaction:** \( 2 \text{O}_3 \rightarrow 3 \text{O}_2 \)

- **Oxidation States:**
- Ozone (\( \text{O}_3 \)): Each oxygen has an oxidation state of 0.
- Oxygen (\( \text{O}_2 \)): Each oxygen also has an oxidation state of 0.

- **Conclusion:** There is no change in oxidation states. Therefore, this is **not** a disproportionation reaction.

### Step 2: Analyze the Second Reaction
**Reaction:** \( \text{KClO}_3 \rightarrow \text{KCl} + \text{Cl}_2 + \text{O}_2 \)

- **Oxidation States:**
- In \( \text{KClO}_3 \), K is +1, O is -2, and Cl is +5.
- In \( \text{KCl} \), K is +1, Cl is -1.
- In \( \text{Cl}_2 \), Cl is 0.
- In \( \text{O}_2 \), O is 0.

- **Changes:**
- Cl changes from +5 in \( \text{KClO}_3 \) to -1 in \( \text{KCl} \) (reduction).
- Cl changes from +5 in \( \text{KClO}_3 \) to 0 in \( \text{Cl}_2 \) (oxidation).

- **Conclusion:** Since Cl undergoes both oxidation and reduction, this is a **disproportionation reaction**.

### Step 3: Analyze the Third Reaction
**Reaction:** \( 2 \text{H}_2\text{O}_2 \rightarrow 2 \text{H}_2\text{O} + \text{O}_2 \)

- **Oxidation States:**
- In \( \text{H}_2\text{O}_2 \), H is +1, O is -1.
- In \( \text{H}_2\text{O} \), H is +1, O is -2.
- In \( \text{O}_2 \), O is 0.

- **Changes:**
- O changes from -1 in \( \text{H}_2\text{O}_2 \) to -2 in \( \text{H}_2\text{O} \) (reduction).
- O changes from -1 in \( \text{H}_2\text{O}_2 \) to 0 in \( \text{O}_2 \) (oxidation).

- **Conclusion:** Since O undergoes both oxidation and reduction, this is also a **disproportionation reaction**.

### Step 4: Analyze the Fourth Reaction
**Reaction:** \( 2 \text{K}_2\text{O}_2 \rightarrow 2 \text{K}_2\text{O} + \text{O}_2 \)

- **Oxidation States:**
- In \( \text{K}_2\text{O}_2 \), K is +1, O is -1.
- In \( \text{K}_2\text{O} \), K is +1, O is -2.
- In \( \text{O}_2 \), O is 0.

- **Changes:**
- O changes from -1 in \( \text{K}_2\text{O}_2 \) to -2 in \( \text{K}_2\text{O} \) (reduction).
- O changes from -1 in \( \text{K}_2\text{O}_2 \) to 0 in \( \text{O}_2 \) (oxidation).

- **Conclusion:** Since O undergoes both oxidation and reduction, this is also a **disproportionation reaction**.

### Final Conclusion
The reactions that are disproportionation reactions are the second, third, and fourth reactions. The first reaction is not a disproportionation reaction.

### Summary of Results
- **Disproportionation Reactions:** 2nd, 3rd, and 4th reactions.
- **Not a Disproportionation Reaction:** 1st reaction.