How many Faradays of electricity are required for the given reaction to occur?
`MnO_(4)^(-) rarr Mn^(2+)`

Options

A. 5F
B. 3F
C. 1F
D. 7F

Answer

To determine how many Faradays of electricity are required for the reaction where \( \text{MnO}_4^- \) is converted to \( \text{Mn}^{2+} \), we can follow these steps:

### Step 1: Identify the oxidation states
In the reaction, manganese in \( \text{MnO}_4^- \) has an oxidation state of +7, and in \( \text{Mn}^{2+} \), it has an oxidation state of +2.

### Step 2: Calculate the change in oxidation state
The change in oxidation state from +7 to +2 indicates a reduction. The total change in oxidation state is:
\[ +7 \text{ (in MnO}_4^-) - (+2 \text{ (in Mn}^{2+}) = 5 \]
This means that manganese is reduced by 5 units.

### Step 3: Determine the number of electrons required
Since each unit of change in oxidation state corresponds to the transfer of one electron, a change of 5 units means that 5 electrons are required for the reduction of one mole of \( \text{MnO}_4^- \) to \( \text{Mn}^{2+} \).

### Step 4: Relate electrons to Faradays
1 Faraday (F) is the charge of one mole of electrons, which is approximately 96500 coulombs. Therefore, if 5 moles of electrons are needed, the total charge required is:
\[ 5 \text{ moles of electrons} = 5 \text{ Faradays} \]

### Conclusion
Thus, the number of Faradays of electricity required for the reaction \( \text{MnO}_4^- \rightarrow \text{Mn}^{2+} \) is **5 Faradays**.