Hydrogen peroxide (H₂O₂) can act as both an oxidizing agent and a reducing agent, depending on the reactants and the conditions of the reaction. Let's illustrate both properties using chemical equations involving first row transition metal ions.

Hydrogen peroxide as an oxidizing agent:

In an alkaline solution, hydrogen peroxide can oxidize manganese(II) ions (Mn²⁺) to manganese dioxide (MnO₂). The balanced chemical equation for this reaction is:

$$2\text{Mn}^{2+} + \text{H}_2\text{O}_2 + 4\text{OH}^- \rightarrow 2\text{MnO}_2 + 4\text{H}_2\text{O}$$

In this reaction, H₂O₂ is reduced to water (H₂O), while Mn²⁺ is oxidized to MnO₂.

Hydrogen peroxide as a reducing agent:

Conversely, in an alkaline solution, hydrogen peroxide can reduce chromium(VI) ions (CrO₄^2-) to chromium(III) ions (Cr³⁺). The balanced chemical equation for this reaction is:

$$2\text{CrO}_4^{2-} + 3\text{H}_2\text{O}_2 + 2\text{OH}^- \rightarrow 2\text{Cr}^{3+} + 3\text{O}_2 + 5\text{H}_2\text{O}$$

In this reaction, H₂O₂ is oxidized to oxygen (O₂), while CrO₄^2- is reduced to Cr³⁺.

Through these equations, it is clear that hydrogen peroxide can exhibit both oxidizing and reducing properties in alkaline solutions, depending on the nature of the transition metal ion it interacts with. These dual roles make H₂O₂ a versatile reagent in various chemical processes.