Summary of "balancing a redox reaction / oxidation number method"

Overview

The video demonstrates balancing the redox reaction between permanganate (MnO4–) and iron(II) (Fe2+) in acidic solution using the oxidation-number method. Water (H2O) and hydrogen ions (H+) are used to balance O and H because the reaction occurs in acid.

Step-by-step method

Assign oxidation numbers to the species that change oxidation state:
- Mn in MnO4–: +7
- Fe in Fe2+: +2
- Mn2+ (product): +2
- Fe3+ (product): +3
- Note: for monatomic ions (Mn2+, Fe2+, Fe3+) the oxidation state equals the ionic charge.
Write the two changing species as paired half-reactions (for comparison):
- MnO4– → Mn2+
- Fe2+ → Fe3+
Determine the electron changes by taking the difference in oxidation numbers:
- Mn: +7 → +2 is a gain of 5 electrons (reduction).
- Fe: +2 → +3 is a loss of 1 electron (oxidation).
Identify oxidation/reduction and label agents:
- MnO4– (permanganate) is the oxidizing agent (it is reduced).
- Fe2+ is the reducing agent (it is oxidized).
Scale the half-reactions so electrons lost = electrons gained:
- Multiply the Fe2+ → Fe3+ half-reaction by 5 so both involve 5 electrons.
Balance oxygen and hydrogen (acidic solution):
- Balance O by adding H2O: MnO4– has 4 O → add 4 H2O on the product side.
- Balance H by adding H+: 4 H2O has 8 H → add 8 H+ to the reactant side.
Combine and check:
- Verify atom counts on each side.
- Verify total charge on each side (must match). In this example the total charge on both sides is +17, confirming the balance.

Balanced net ionic equation

MnO4– + 5 Fe2+ + 8 H+ → Mn2+ + 5 Fe3+ + 4 H2O

Key concepts emphasized

Use oxidation numbers to determine how many electrons are transferred.
Scale half-reactions so electrons lost = electrons gained.
In acidic medium, use H2O to balance oxygen and H+ to balance hydrogen.
Always perform final atom and charge checks to confirm the equation is balanced.
Identify oxidizing and reducing agents by which species gain or lose electrons.