To determine in which of the given reactions `H2PO4^(-)` acts as an acid, we need to analyze each reaction based on the definitions of acids according to different theories (Arrhenius, Brønsted-Lowry, and Lewis).
### Step-by-Step Solution:
1. **Understanding Acids**:
- According to the Arrhenius theory, an acid is a substance that increases the concentration of H⁺ ions in aqueous solution.
- According to the Brønsted-Lowry theory, an acid is a proton donor.
- According to the Lewis theory, an acid is an electron pair acceptor.
2. **Analyzing Reaction I**:
- The reaction is:
\[ H3PO4 + H2O \rightleftharpoons H3O^+ + H2PO4^- \]
- In this reaction, `H3PO4` donates a proton (H⁺) to water, forming `H3O^+` and `H2PO4^-`.
- Here, `H3PO4` acts as an acid, while `H2PO4^-` is formed as a product and does not act as an acid in this reaction.
3. **Analyzing Reaction II**:
- The reaction is:
\[ H2PO4^- + H2O \rightleftharpoons HPO4^{2-} + H3O^+ \]
- In this case, `H2PO4^-` donates a proton (H⁺) to water, resulting in the formation of `HPO4^{2-}` and `H3O^+`.
- Here, `H2PO4^-` acts as an acid because it donates a proton.
4. **Analyzing Reaction III**:
- The reaction is:
\[ H2PO4^- + OH^- \rightleftharpoons H3PO4 + O^{2-} \]
- In this reaction, `H2PO4^-` accepts a proton from `OH^-`, resulting in the formation of `H3PO4` and `O^{2-}`.
- Here, `H2PO4^-` acts as a base because it accepts a proton.
5. **Conclusion**:
- From the analysis, `H2PO4^-` acts as an acid only in Reaction II. In Reaction I, it does not act as an acid, and in Reaction III, it acts as a base.
### Final Answer:
`H2PO4^-` acts as an acid only in Reaction II.
