For the electrolytic production of `NaClO_(4)` from `NaClO_(3)` according to the reaction `NaClO_(3)+H_(2)O rarr NaClO_(4)+H_(2). ` How many faradays of electricity would be required to produce `0.5 mo l e` of `NaClO_(4)`?

Options

A. 1
B. 2
C. 3
D. `1.5`

Explanation

To solve the problem of how many faradays of electricity are required to produce 0.5 moles of sodium perchlorate (NaClO₄) from sodium chlorate (NaClO₃), we can follow these steps:

### Step-by-Step Solution:

1. **Write the Given Reaction:**
The reaction for the electrolytic production of sodium perchlorate from sodium chlorate is:
\[ \text{NaClO}_3 + \text{H}_2\text{O} \rightarrow \text{NaClO}_4 + \text{H}_2 \]

2. **Identify the Change in Oxidation State:**
In this reaction, the chlorate ion (ClO₃⁻) is being converted to the perchlorate ion (ClO₄⁻). The oxidation state of chlorine changes from +5 in ClO₃⁻ to +7 in ClO₄⁻.

3. **Determine the Number of Electrons Involved:**
The conversion of ClO₃⁻ to ClO₄⁻ involves the loss of 2 electrons:
\[ \text{ClO}_3^- + 2\text{H}^+ + 2e^- \rightarrow \text{ClO}_4^- + \text{H}_2 \]
This shows that 2 moles of electrons are required to produce 1 mole of ClO₄⁻.

4. **Calculate the Faradays Required for 0.5 Moles of NaClO₄:**
Since 2 moles of electrons (or 2 faradays) are required to produce 1 mole of ClO₄⁻, for 0.5 moles of ClO₄⁻, the amount of electricity required will be:
\[ \text{Faradays required} = 0.5 \text{ moles ClO}_4^- \times \frac{2 \text{ faradays}}{1 \text{ mole ClO}_4^-} = 1 \text{ faraday} \]

5. **Conclusion:**
Therefore, to produce 0.5 moles of sodium perchlorate (NaClO₄), **1 faraday** of electricity is required.

### Final Answer:
**1 Faraday**