Which of the following is the correctly balanced equation for the reaction of MnO4^- with Fe^2+ in acidic conditions?

• A. 5 Fe^{2+} + MnO_4^- + 8 H^+ → 5 Fe^{3+} + Mn^{2+} + 4 H_2O
• B. 4 Fe^{2+} + MnO_4^- + 10 H^+ → 4 Fe^{3+} + Mn^{2+} + 5 H_2O
• C. 3 Fe^{2+} + MnO_4^- + 6 H^+ → 3 Fe^{3+} + Mn^{2+} + 3 H_2O
• D. 2 Fe^{2+} + MnO_4^- + 4 H^+ → 2 Fe^{3+} + Mn^{2+} + 2 H_2O

Answer: (A)

Verified step by step guidance

1. Identify the oxidation and reduction half-reactions in the given redox reaction. Manganese in MnO_4^- is reduced from +7 to +2 in Mn^{2+}, and iron Fe^{2+} is oxidized to Fe^{3+}.
2. Write the half-reactions separately: the reduction half-reaction for MnO_4^- to Mn^{2+} and the oxidation half-reaction for Fe^{2+} to Fe^{3+}.
3. Balance each half-reaction for atoms other than H and O first, then balance oxygen atoms by adding H_2O, and balance hydrogen atoms by adding H^+ ions (since the reaction is in acidic conditions).
4. Balance the charges in each half-reaction by adding electrons (e^-) to the appropriate side so that the total charge is the same on both sides of each half-reaction.
5. Multiply the half-reactions by appropriate coefficients so that the number of electrons lost in oxidation equals the number gained in reduction, then add the half-reactions together, canceling electrons and any species that appear on both sides to get the balanced overall equation.