How many coulombs are required for the oxidation of 1 mol of FeO to `Fe_(2)O_(3)` ?
(Hint. `Fe^(2+) to Fe^(3+)+e^(-)`)

Explanation

To determine how many coulombs are required for the oxidation of 1 mol of FeO to Fe₂O₃, we can follow these steps:

### Step 1: Identify the oxidation states
In FeO, iron (Fe) is in the +2 oxidation state (Fe²⁺), and in Fe₂O₃, iron is in the +3 oxidation state (Fe³⁺).

### Step 2: Write the half-reaction for oxidation
The oxidation process can be represented as:
\[ \text{Fe}^{2+} \rightarrow \text{Fe}^{3+} + e^- \]
This indicates that one electron is lost during the oxidation of one Fe²⁺ ion to Fe³⁺.

### Step 3: Determine the number of moles of Fe involved
In the conversion of 1 mol of FeO to Fe₂O₃, we have 2 moles of Fe in Fe₂O₃. Therefore, we need to consider the oxidation of 2 moles of Fe²⁺ ions (from 2 moles of FeO) to 2 moles of Fe³⁺ ions.

### Step 4: Calculate the total number of electrons transferred
Since each Fe²⁺ ion loses 1 electron, for 2 moles of Fe²⁺, the total number of electrons transferred is:
\[ 2 \text{ moles of Fe}^{2+} \times 1 \text{ electron/mole} = 2 \text{ electrons} \]

### Step 5: Use Faraday's constant to find the total charge
Faraday's constant (F) is approximately 96500 coulombs per mole of electrons. Therefore, for 2 moles of electrons, the total charge (Q) can be calculated as:
\[ Q = n \times F \]
Where \( n \) is the number of moles of electrons. Thus:
\[ Q = 2 \, \text{moles} \times 96500 \, \text{C/mole} = 193000 \, \text{C} \]

### Conclusion
The total charge required for the oxidation of 1 mol of FeO to Fe₂O₃ is **193000 coulombs**.

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