Which of the following solutions will exactly oxidize `25mL` of an acid solution of `0.1 M Fe` (`II`) oxalate?

Options

A. `25 mL` of `0.1 M KMnO_(4)`
B. `25 mL` of `0.2 M KMnO_(4)`
C. `25 mL` of `0.6 M KMnO_(4)`
D. `15 mL` of `0.1 M KMnO_(4)`

Explanation

To determine which solution will exactly oxidize 25 mL of a 0.1 M Fe(II) oxalate solution, we need to follow these steps:

### Step 1: Understand the Reaction
The oxidation of Fe(II) oxalate (FeC2O4) involves the reduction of permanganate ions (MnO4^-). The balanced reaction is:
\[ 5 \text{Fe}^{2+} + 3 \text{MnO}_4^- + 6 \text{H}^+ \rightarrow 5 \text{Fe}^{3+} + 3 \text{Mn}^{2+} + 3 \text{CO}_2 + 3 \text{H}_2O \]

### Step 2: Calculate Moles of Fe(II) Oxalate
We need to calculate the moles of Fe(II) in the solution:
- Volume of Fe(II) solution = 25 mL = 0.025 L
- Concentration of Fe(II) = 0.1 M

Using the formula:
\[ \text{Moles of Fe}^{2+} = \text{Concentration} \times \text{Volume} \]
\[ \text{Moles of Fe}^{2+} = 0.1 \, \text{mol/L} \times 0.025 \, \text{L} = 0.0025 \, \text{mol} \]

### Step 3: Determine Moles of MnO4^- Required
From the balanced equation, 5 moles of Fe(II) react with 3 moles of MnO4^-:
- Moles of MnO4^- required = (3/5) × Moles of Fe(II)
\[ \text{Moles of MnO4^-} = \frac{3}{5} \times 0.0025 \, \text{mol} = 0.0015 \, \text{mol} \]

### Step 4: Calculate the Volume of MnO4^- Solution Required
To find the volume of the MnO4^- solution needed, we need to know its concentration. Assuming we have different options for concentrations, we can calculate the volume required for each concentration.

Using the formula:
\[ \text{Volume} = \frac{\text{Moles}}{\text{Concentration}} \]

For each concentration option (let's denote them as M1, M2, M3, etc.), the volume can be calculated as:
\[ \text{Volume of MnO4^-} = \frac{0.0015 \, \text{mol}}{M} \]

### Step 5: Compare with Given Options
Now, calculate the volume for each concentration option provided in the question. For example:
1. If M1 = 0.1 M:
\[ \text{Volume} = \frac{0.0015}{0.1} = 0.015 \, \text{L} = 15 \, \text{mL} \]
2. If M2 = 0.2 M:
\[ \text{Volume} = \frac{0.0015}{0.2} = 0.0075 \, \text{L} = 7.5 \, \text{mL} \]
3. If M3 = 0.6 M:
\[ \text{Volume} = \frac{0.0015}{0.6} = 0.0025 \, \text{L} = 2.5 \, \text{mL} \]

### Conclusion
You will need to calculate the volume for each concentration option provided in the question to find out which one will exactly oxidize the 25 mL of 0.1 M Fe(II) oxalate solution.