Consider the cell potentials `E_(Mg^(2+)|Mg)^(0)=-2.37V` and `E_(Fe^(3+)|Fe)^(0)=-0.04V`
The best reducing agent would be

Options

A. `Mg^(2+)`
B. `Fe^(3+)` decrease
C. `Mg`
D. `Fe`

Explanation

To determine the best reducing agent from the given cell potentials, we can follow these steps:

### Step 1: Understand the Given Potentials
We have two half-cell potentials:
- \( E^\circ_{Mg^{2+}/Mg} = -2.37 \, V \)
- \( E^\circ_{Fe^{3+}/Fe} = -0.04 \, V \)

### Step 2: Convert Reduction Potentials to Oxidation Potentials
To find the oxidation potentials, we need to change the sign of the reduction potentials:
- For magnesium:
\[ E^\circ_{Mg/Mg^{2+}} = -E^\circ_{Mg^{2+}/Mg} = -(-2.37) = 2.37 \, V \]
- For iron:
\[ E^\circ_{Fe/Fe^{2+}} = -E^\circ_{Fe^{3+}/Fe} = -(-0.04) = 0.04 \, V \]

### Step 3: Compare the Oxidation Potentials
Now we compare the oxidation potentials:
- \( E^\circ_{Mg/Mg^{2+}} = 2.37 \, V \)
- \( E^\circ_{Fe/Fe^{2+}} = 0.04 \, V \)

Since \( 2.37 \, V > 0.04 \, V \), magnesium has a higher oxidation potential than iron.

### Step 4: Identify the Strongest Reducing Agent
A reducing agent is a substance that donates electrons and gets oxidized in the process. The stronger the reducing agent, the higher its oxidation potential. Since magnesium has a higher oxidation potential, it is the stronger reducing agent compared to iron.

### Conclusion
Thus, the best reducing agent among the given options is:
**Magnesium (Mg)**

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