Using the standard electrode potential values given below, decide which of the statements, `I,II,III` and `IV` are correct. Choose the right answer from `(1) (2)` and `(4)`
`Fe^(2+)+2e^(-) hArr " "," "E^(0)=-0.44V`
`Cu^(2+)+2e^(-) hArr Cu" "," "E^(0)=+0.34V`
`Ag^(+)+e^(-)hArr Fe" "," "E^(0)=+0.80V`
`I. ` Copper can displace iron from `FeSO_(4)` solution.
`II.` Iron can displace copper from `CuSO_(4)` solution
`III.` Silver can displace copper from `CuSO_(4)` solution
`IV.` Iron can displace silver from `AgNO_(3)` solution.

Options

A. I and II
B. II and III
C. II and IV
D. I and IV

Answer

To solve the problem, we will analyze each statement based on the standard electrode potentials provided. The key is to determine whether a metal can displace another metal from its salt solution based on their electrode potentials.

### Given Standard Electrode Potentials:
1. \( \text{Fe}^{2+} + 2e^- \leftrightarrow \text{Fe} \) \( E^{0} = -0.44 \, \text{V} \)
2. \( \text{Cu}^{2+} + 2e^- \leftrightarrow \text{Cu} \) \( E^{0} = +0.34 \, \text{V} \)
3. \( \text{Ag}^{+} + e^- \leftrightarrow \text{Ag} \) \( E^{0} = +0.80 \, \text{V} \)

### Analyzing Each Statement:

**Statement I: Copper can displace iron from FeSO₄ solution.**

- The reaction can be written as:
\[ \text{Cu} + \text{Fe}^{2+} \rightarrow \text{Cu}^{2+} + \text{Fe} \]
- Calculate \( E^{0}_{cell} \):
\[ E^{0}_{cell} = E^{0}_{cathode} - E^{0}_{anode} = (-0.44) - (+0.34) = -0.78 \, \text{V} \]
- Since \( E^{0}_{cell} < 0 \), this reaction is **not possible**.

**Statement II: Iron can displace copper from CuSO₄ solution.**

- The reaction can be written as:
\[ \text{Fe} + \text{Cu}^{2+} \rightarrow \text{Fe}^{2+} + \text{Cu} \]
- Calculate \( E^{0}_{cell} \):
\[ E^{0}_{cell} = E^{0}_{cathode} - E^{0}_{anode} = (+0.34) - (-0.44) = +0.78 \, \text{V} \]
- Since \( E^{0}_{cell} > 0 \), this reaction is **possible**.

**Statement III: Silver can displace copper from CuSO₄ solution.**

- The reaction can be written as:
\[ \text{Ag} + \text{Cu}^{2+} \rightarrow \text{Ag}^{+} + \text{Cu} \]
- Calculate \( E^{0}_{cell} \):
\[ E^{0}_{cell} = E^{0}_{cathode} - E^{0}_{anode} = (+0.34) - (+0.80) = -0.46 \, \text{V} \]
- Since \( E^{0}_{cell} < 0 \), this reaction is **not possible**.

**Statement IV: Iron can displace silver from AgNO₃ solution.**

- The reaction can be written as:
\[ \text{Fe} + \text{Ag}^{+} \rightarrow \text{Fe}^{2+} + \text{Ag} \]
- Calculate \( E^{0}_{cell} \):
\[ E^{0}_{cell} = E^{0}_{cathode} - E^{0}_{anode} = (+0.80) - (-0.44) = +1.24 \, \text{V} \]
- Since \( E^{0}_{cell} > 0 \), this reaction is **possible**.

### Conclusion:
- Correct statements are **II** and **IV**.
- Therefore, the correct answer is option **(3)**: Statements II and IV are correct.