The reaction represented by the equation: Ni\(^{2+}\)\((_{aq}\)) + Fe\((_s\)) → Ni\((_s\)) + Fe\(^{2+}\)\((_{aq}\)) is a redox reaction because
Ni\(^{2+}\)\(_{aq}\) + Fe\(_s\) → Ni\(_s\) + Fe\(^{2+}\)\(_{aq}\)
Reduction takes place from Ni\(^{2+}\) to Ni because there is a decrease in oxidation number from +2 to 0. Thus, Ni\(^{2+}\) ions are reduced to Ni
Oxidation takes place from Fe to Fe\(^{2+}\) because there is an increase in oxidation number from 0 to +2. Thus, Fe is oxidized to Fe\(^{2+}\)
Oxidizing agent → Ni\(^{2+}\) because oxidizing agents undergo reduction.
Reducing agent → Fe because reducing agents undergo oxidation.
Consequently, option D is the correct option.
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