In the following equation, `MnO_(2)` acts as
`MnO_(4)^(2-) + 2H_(2)O + 2E^(-) rarr MnO_(2) + 4overset(Ө)(OH)`

Options

A. Oxidising agent
B. Reducing agent
C. Both oxidising and reducing agent.
D. Neither oxidising nor reducing agent.

Answer

To determine the role of MnO₂ in the given redox reaction, we need to analyze the oxidation states of manganese in the reactants and products.

### Step-by-Step Solution:

1. **Identify the Reactants and Products:**
The reaction given is:
\[ \text{MnO}_4^{2-} + 2\text{H}_2\text{O} + 2\text{e}^- \rightarrow \text{MnO}_2 + 4\text{OH}^- \]

2. **Determine the Oxidation States:**
- In \(\text{MnO}_4^{2-}\), the oxidation state of manganese (Mn) can be calculated as follows:
\[ \text{Let the oxidation state of Mn be } x. \]
The equation for the oxidation state is:
\[ x + 4(-2) = -2 \implies x - 8 = -2 \implies x = +6. \]
- In \(\text{MnO}_2\), the oxidation state of manganese is:
\[ x + 2(-2) = 0 \implies x - 4 = 0 \implies x = +4. \]

3. **Analyze the Change in Oxidation State:**
- The oxidation state of Mn changes from +6 in \(\text{MnO}_4^{2-}\) to +4 in \(\text{MnO}_2\). This indicates that Mn is being reduced (gaining electrons).

4. **Identify the Role of MnO₂:**
- Since \(\text{MnO}_4^{2-}\) is reduced to \(\text{MnO}_2\), it acts as an oxidizing agent (it causes the oxidation of another species by accepting electrons).
- The species that loses electrons (in this case, water, \(\text{H}_2\text{O}\)) acts as a reducing agent.

5. **Conclusion About MnO₂:**
- In the context of the reaction, \(\text{MnO}_2\) does not undergo any further change in oxidation state and does not act as either an oxidizing or reducing agent. Therefore, it is classified as neither.

### Final Answer:
MnO₂ acts as neither an oxidizing nor a reducing agent in the given reaction.