How many mL of a 0.05 M `KMnO_(4)` solution are required to oxidise 2.0 g of `FeSO_(4)` in a dilute solution (acidic) ?

Explanation

To solve the problem of how many mL of a 0.05 M `KMnO4` solution are required to oxidize 2.0 g of `FeSO4` in a dilute acidic solution, we can follow these steps:

### Step 1: Write the balanced chemical equation
The balanced equation for the reaction between `FeSO4` and `KMnO4` in acidic medium is:
\[ 5 \text{FeSO}_4 + \text{KMnO}_4 + 8 \text{H}_2\text{SO}_4 \rightarrow 5 \text{Fe}_2(\text{SO}_4)_3 + \text{MnSO}_4 + K_2\text{SO}_4 + 8 \text{H}_2\text{O} \]

### Step 2: Calculate the moles of `FeSO4`
To find the moles of `FeSO4`, we use the formula:
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