Which of the below is the correctly balanced redox reaction of ammonium ion (NH4+) to the nitrite ion (NO2−).

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We start by restating the problem and listing the options clearly to compare them.
Question: Which of the below is the correctly balanced redox reaction of ammonium ion (NH4+) to the nitrite ion (NO2−).
Answer options:
1) NH4+(aq) + 2H2O(l) → NO2−(aq) + 8H+(aq) + 6e−
2) NH4+(aq) + H2O(l) → NO2−(aq) + 2H+(aq)
3) NH4+(aq) + 2H2O(l) → NO2−(aq) + 4H+(aq) + 2e−
4) NH4+(aq) + H2O(l) → NO2−(aq) + H+(aq)
5) NH4+(aq) + 2H2O(l) → NO2−(aq) + 8H+(aq) + 2e−

Analyzing each option:
Option 1: NH4+(aq) + 2H2O(l) → NO2−(aq) + 8H+(aq) + 6e−
- Check oxidation states: N in NH4+ is −3; in NO2− it is +3. Change is +6 electrons lost, so 6e− should be produced. This aligns with the oxidation of N from −3 to +3.
- Balance of atoms: N: 1 on both sides. O: 2 on both sides (NO2− has 2 O; reactants include 2 H2O giving 2 O total). H: 4 H in NH4+ plus 4 H from 2H2O equals 8 H on the left; 8H+ on the right provides 8 H. Charge: Left has +1 (NH4+); right has NO2− (−1) + 8H+ (+8) + 6e− (−6) = +1. Overall, all elements and charges balance. This option is internally consistent and correctly balanced in acidic solution.
Option 2: NH4+(aq) + H2O(l) → NO2−(aq) + 2H+(aq)
- Atom balance: Left has 1 N and 2 H2O contributing additional H and O; right has NO2− (2 O) but only 2 H+. Oxygen balance is likely off because NO2− contains 2 O but the left side provides only 1 O from H2O, so O is not balanced. Also electron balance is not shown and would be inconsistent with a proper redox half-reaction. This option is not a correctly balanced equation.
Option 3: NH4+(aq) + 2H2O(l) → NO2−(aq) + 4H+(aq) + 2e−
- Here the electrons produced are 2e−, which would imply only a −2 change in oxidation state, contradicting the actual −3 to +3 change (which requires 6 electrons). This makes the electron balance incorrect, so the overall redox balance is invalid.
Option 4: NH4+(aq) + H2O(l) → NO2−(aq) + H+(aq)
- This option shows 1 H+ on the product side, implying partial balancing of hydrogen. Oxygen balance is also problematic since NO2− contains 2 oxygens but only one oxygen source is present on the left (H2O). Electron balance is not addressed, so this cannot be a proper balanced redox equation.
Option 5: NH4+(aq) + 2H2O(l) → NO2−(aq) + 8H+(aq) + 2e−
- While hydrogen and nitrogen atoms appear balanced, the electron count is only 2e−, which would not account for the actual oxidation change of −3 to +3 (requires 6 electrons). This mismatch makes the half-reaction incorrect.

Summary of why Option 1 is correct and the others are not: the correct balanced half-reaction must conserve atoms (N, O, H) and satisfy charge balance, along with producing the correct number of electrons corresponding to the oxidation state change of nitrogen from −3 in NH4+ to +3 in NO2−. Option 1 satisfies all these criteria: N balance (1 on both sides), O balance (2 O on both sides via 2H2O on the left and NO2− on the right), H balance (8 H on the left from NH4+ and 2H2O, matched by 8H+ on the right), and charge balance (+1 on the left, +1 on the right when accounting for NO2−, 8H+, and 6e−). The remaining options fail due to incorrect atom balance, incorrect electron balance, or incomplete charge balance, revealing fundamental errors in redox accounting or in the stoichiometry.

AGRI10046_2025_SM2 AGRI10046 PRACTICE exam

Which of the below is the correctly balanced redox reaction of ammonium ion (NH4+) to the nitrite ion (NO2−).

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