The number of electrons involved when 1 mole of `H_(2)O_(2)` decomposes as
`H_(2)O_(2) rarr H_(2)O +O_(2)`, is

Options

A. 1
B. 8
C. 6
D. 4

Answer

To determine the number of electrons involved when 1 mole of \( H_2O_2 \) decomposes into \( H_2O \) and \( O_2 \), we can follow these steps:

### Step 1: Write the balanced equation
The decomposition of hydrogen peroxide can be represented by the equation:
\[ H_2O_2 \rightarrow H_2O + O_2 \]

### Step 2: Determine oxidation states
Next, we need to assign oxidation states to each element in the reaction:
- In \( H_2O_2 \) (hydrogen peroxide), the oxidation state of oxygen is -1.
- In \( H_2O \) (water), the oxidation state of oxygen is -2.
- In \( O_2 \) (molecular oxygen), the oxidation state of oxygen is 0.

### Step 3: Identify changes in oxidation states
Now, we can identify the changes in oxidation states:
- For the oxygen in \( H_2O_2 \) going to \( H_2O \):
- Change from -1 to -2 (reduction).
- For the oxygen in \( H_2O_2 \) going to \( O_2 \):
- Change from -1 to 0 (oxidation).

### Step 4: Calculate the number of electrons transferred
Each oxygen atom in \( H_2O_2 \) undergoes a change in oxidation state:
- The reduction (from -1 to -2) involves 1 electron per oxygen atom.
- The oxidation (from -1 to 0) involves 1 electron per oxygen atom.

Since there are 2 oxygen atoms in \( H_2O_2 \):
- The total number of electrons involved in reduction: \( 1 \text{ electron} \times 1 \text{ atom} = 1 \text{ electron} \)
- The total number of electrons involved in oxidation: \( 1 \text{ electron} \times 1 \text{ atom} = 1 \text{ electron} \)

### Step 5: Total number of electrons
Adding these together, the total number of electrons transferred during the decomposition of 1 mole of \( H_2O_2 \) is:
\[ 1 \text{ (from reduction)} + 1 \text{ (from oxidation)} = 2 \text{ electrons} \]

### Conclusion
Thus, the number of electrons involved when 1 mole of \( H_2O_2 \) decomposes is **2 electrons**.

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