If we balanced the following reaction, what is the net ionic charge on the right side of the equation? H+ + MnO4- + Fe2+ →Mn2+ + Fe3+ + H2O
Correct answer: D. +17
- A. +5
- B. +7
- C. +10
- D. +17
Explanation
To solve this problem, we must balance the redox reaction by ensuring that both mass and charge are conserved. Manganese is reduced from +7 in MnO4- to +2 in Mn2+, gaining 5 electrons. Iron is oxidized from +2 in Fe2+ to +3 in Fe3+, losing 1 electron per iron atom. By balancing the electron transfer, we multiply the iron half-equation by 5, resulting in the balanced equation:8H+ + MnO4- + 5Fe2+ → Mn2+ + 5Fe3+ + 4H2OOn the right side, the charges are:Mn2+: +25Fe3+: +15 (5 x +3)H2O: NeutralTotal charge on the right: +2 + 15 = +17Thus, the correct net ionic charge on the right side is +17. Other options are incorrect as they do not reflect the correct balance of charges based on the stoichiometry of the reaction.
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About Balancing Redox Equations
Redox equations are balanced by tracking electron loss and gain through oxidation numbers or the ion-electron method. Work includes identifying oxidizing and reducing agents, balancing atoms and charge in acidic or basic medium, and ensuring that electrons lost equal electrons gained.
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