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

• A. 2 Fe^2+ + MnO4^- + 4 H^+ → 2 Fe^3+ + MnO2 + 2 H2O
• B. 4 Fe^2+ + MnO4^- + 8 H^+ → 4 Fe^3+ + MnO2 + 4 H2O
• C. 5 Fe^2+ + MnO4^- + 8 H^+ → 5 Fe^3+ + Mn^2+ + 4 H2O
• D. 3 Fe^2+ + MnO4^- + 6 H^+ → 3 Fe^3+ + MnO2 + 3 H2O

Answer: (C)

Verified step by step guidance

1. Identify the oxidation states of manganese and iron in the reactants and products to determine which species are oxidized and which are reduced. Mn in MnO4^- is typically +7, Fe^2+ is +2, and in the products, Fe^3+ is +3, MnO2 contains Mn in +4, and Mn^2+ is +2.
2. Write the half-reactions for oxidation and reduction separately. For example, the oxidation half-reaction involves Fe^2+ being oxidized to Fe^3+, and the reduction half-reaction involves MnO4^- being reduced to either MnO2 or Mn^2+, depending on the product given.
3. Balance each half-reaction for atoms other than H and O first, then balance oxygen atoms by adding H2O, and balance hydrogen atoms by adding H^+ ions since the solution is acidic.
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 the half-reaction.
5. Combine the two half-reactions by multiplying them by appropriate coefficients so that the electrons cancel out, then add the half-reactions together and simplify to get the balanced overall equation.