Which of the following is the correctly balanced redox reaction occurring in a basic solution for the given reactants and products: ClO⁻(aq) + Cr(OH)₄⁻(aq) → CrO₄²⁻(aq) + Cl⁻(aq)?

• A. ClO⁻(aq) + Cr(OH)₄⁻(aq) → CrO₄²⁻(aq) + Cl⁻(aq)
• B. 3 ClO⁻(aq) + 2 Cr(OH)₄⁻(aq) + 4 OH⁻(aq) → 2 CrO₄²⁻(aq) + 3 Cl⁻(aq) + 4 H₂O(l)
• C. ClO⁻(aq) + 2 Cr(OH)₄⁻(aq) + OH⁻(aq) → 2 CrO₄²⁻(aq) + Cl⁻(aq) + 2 H₂O(l)
• D. 2 ClO⁻(aq) + Cr(OH)₄⁻(aq) + 2 OH⁻(aq) → CrO₄²⁻(aq) + 2 Cl⁻(aq) + H₂O(l)

Answer: (B)

Verified step by step guidance

1. Identify the oxidation and reduction half-reactions. In this case, ClO⁻ is reduced to Cl⁻, and Cr(OH)₄⁻ is oxidized to CrO₄²⁻.
2. Balance each half-reaction separately. For the reduction half-reaction, balance ClO⁻ to Cl⁻ by adding electrons. For the oxidation half-reaction, balance Cr(OH)₄⁻ to CrO₄²⁻ by adding electrons and water molecules.
3. Since the reaction occurs in a basic solution, add OH⁻ ions to both sides of each half-reaction to neutralize any H⁺ ions and form water molecules.
4. Combine the balanced half-reactions, ensuring that the number of electrons lost in the oxidation half-reaction equals the number of electrons gained in the reduction half-reaction.
5. Verify that the final equation is balanced in terms of both mass and charge, ensuring that the number of atoms of each element and the total charge are the same on both sides of the equation.