`E^@` values of some redox couples are given below . On the basis of these values choose the correct option .
`E^@` value : `Br_2|Br^(-) =+1.90,Ag^(+)|Ag (s) = +0.80`
`Cu^(2+)|Cu(s)=+0.34, I_2(s)|I^(-)=+0.54`

Options

A. Cu will reduce `Br^(-)`
B. Cu will reduce Ag
C. Cu will reduce `I^-`
D. Cu will reduce `Br_2`

Answer

To solve the question regarding the redox couples and their standard reduction potentials (E° values), we need to determine which species can be reduced by another based on their E° values. The higher the E° value, the stronger the oxidizing agent, and thus the more likely it is to be reduced.

### Given E° values:
1. \( \text{Br}_2 | \text{Br}^- = +1.90 \)
2. \( \text{Ag}^+ | \text{Ag (s)} = +0.80 \)
3. \( \text{Cu}^{2+} | \text{Cu (s)} = +0.34 \)
4. \( \text{I}_2 (s) | \text{I}^- = +0.54 \)

### Step-by-step Solution:

1. **Identify the Reducing and Oxidizing Agents**:
- A species with a higher E° value can reduce a species with a lower E° value.
- Thus, we will compare the E° values to determine which species can reduce which.

2. **Analyze Each Option**:
- **Option A: Can Cu reduce Br2?**
- E° of \( \text{Br}_2 \) is +1.90 (higher than Cu's +0.34).
- Therefore, Cu cannot reduce Br2.

- **Option B: Can Cu reduce Ag?**
- E° of \( \text{Ag}^+ \) is +0.80 (higher than Cu's +0.34).
- Therefore, Cu cannot reduce Ag.

- **Option C: Can Cu reduce I-?**
- Here, we need to consider the reverse reaction: \( \text{I}_2 \) to \( \text{I}^- \) with E° = +0.54.
- Since +0.54 is higher than +0.34, Cu cannot reduce I-.

- **Option D: Can Cu reduce Br-?**
- This option is a bit tricky. We need to consider the reverse of the reaction \( \text{Br}^- \) to \( \text{Br}_2 \).
- The E° for \( \text{Br}_2 \) is +1.90, which is much higher than Cu's +0.34.
- Therefore, Cu can reduce Br2, as it can donate electrons to Br2 and be oxidized.

3. **Conclusion**:
- The only correct option is that Cu can reduce \( \text{Br}_2 \).

### Final Answer:
- **Correct Option: D - Cu can reduce Br2.**