Consider the following unbalanced redox equation: H+ + BrO3– + Fe2+ → Fe3+ + Br– + H2O When this equation is completely balanced using simplest whole number ratio, the number of electrons transferred will be

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Consider the following unbalanced redox equation: H+ +  BrO3– +  Fe2+ →  Fe3+ + Br– + H2O When this equation is completely balanced using simplest whole number ratio, the number of electrons transferred will be

BlackTom AI · Accepted Answer

To tackle this redox equation, I’ll break down the oxidation and reduction half-reactions and then balance the electrons.
First, identify the oxidation states: Fe in Fe2+ goes to Fe3+, so iron is oxidized, losing 1 electron per Fe2+ → Fe3+.
Next, bromine in BrO3− goes to Br−. In BrO3−, Br must have an oxidation state of +5. In Br−, Br is −1. Thus Br changes by gaining 6 electrons (from +5 to −1), i.e., BrO3− + 6 H+ + 6 e− → Br− + 3 H2O is the reduction half-reaction.
Now compare electron bookkeeping: the bromine reduction half consumes 6 electrons per bromate ion reduced, while the iron oxidation half produces 1 electron per Fe2+ oxidized. To balance the electrons in the overall reaction, 6 Fe2+ ions must be oxidized to supply 6 electrons for the BrO3− reduction.
Therefore, the total number of electrons transferred in the balanced overall equation is 6 electrons. This matches the requirement that the electrons produced equal those consumed, with the simplest whole-number ratio.
In summary, each BrO3− requires 6 electrons to be reduced to Br−, and 6 Fe2+ ions supply those electrons by being oxidized to Fe3+. Hence the correct choice is 6.

MUF0042 Chemistry Unit 2 - Semester 2, 2025 10.13 Quiz: Balancing Redox Reactions

Consider the following unbalanced redox equation: H+ + BrO3– + Fe2+ → Fe3+ + Br– + H2O When this equation is completely balanced using simplest whole number ratio, the number of electrons transferred will be

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