Question

Question on balancing redox reaction and their workings

Answer 1

[AI-generated answer — Bot GPT 3.5]

To balance a redox reaction, you need to follow a few steps:

1. Assign oxidation numbers to all the elements in the reaction. This will help you determine which elements are being oxidized (increasing in oxidation state) and which are being reduced (decreasing in oxidation state). Remember that the sum of oxidation numbers in a compound should be zero, and the sum of oxidation numbers in a polyatomic ion should be equal to the charge of the ion.

2. Identify the substance that is being oxidized and write the half-reaction for the oxidation process. The species being oxidized will lose electrons, so its oxidation half-reaction will show an increase in oxidation state and have electrons as a reactant.

3. Identify the substance that is being reduced and write the half-reaction for the reduction process. The species being reduced will gain electrons, so its reduction half-reaction will show a decrease in oxidation state and have electrons as a product.

4. Balance the atoms in each half-reaction. This involves adjusting coefficients so that the number of atoms of each element is the same on both sides of the half-reaction.

5. Balance the charges in each half-reaction. Add electrons to one side of the half-reaction to ensure that the total charge is the same on both sides.

6. Multiply the half-reactions by appropriate coefficients to ensure that the number of electrons gained in the reduction half-reaction is equal to the number of electrons lost in the oxidation half-reaction. This will allow the electrons to cancel out when the half-reactions are added together.

7. Add the balanced half-reactions together to form the overall balanced equation.

For example, let's balance the redox reaction:

Fe2+(aq) + MnO4-(aq) → Fe3+(aq) + Mn2+(aq)

1. Assign oxidation numbers:
Fe2+: +2
MnO4-: Mn +7, O -2 (since the sum of oxidation numbers in a polyatomic ion is equal to its charge)
Fe3+: +3
Mn2+: +2

2. Identify the oxidation half-reaction:
Fe2+(aq) → Fe3+(aq)
Fe2+ is being oxidized, so its half-reaction is:
Fe2+(aq) → Fe3+(aq) + e-

3. Identify the reduction half-reaction:
MnO4-(aq) → Mn2+(aq)
MnO4- is being reduced, so its half-reaction is:
MnO4-(aq) + 8 H+(aq) + 5 e- → Mn2+(aq) + 4 H2O(l)

4. Balance the atoms in each half-reaction:
Oxidation half-reaction: Fe2+(aq) → Fe3+(aq) + e-
Reduction half-reaction: MnO4-(aq) + 8 H+(aq) + 5 e- → Mn2+(aq) + 4 H2O(l)

5. Balance the charges in each half-reaction:
Oxidation half-reaction: Fe2+(aq) → Fe3+(aq) + e-
Reduction half-reaction: MnO4-(aq) + 8 H+(aq) + 5 e- → Mn^2+(aq) + 4 H2O(l) + 5 e-

6. Multiply the half-reactions to balance the electrons:
Oxidation half-reaction: 5 Fe2+(aq) → 5 Fe3+(aq) + 5 e-
Reduction half-reaction: MnO4-(aq) + 8 H+(aq) + 5 e- → Mn2+(aq) + 4 H2O(l) + 5 e-

7. Add the balanced half-reactions together:
5 Fe2+(aq) + MnO4-(aq) + 8 H+(aq) → 5 Fe3+(aq) + Mn2+(aq) + 4 H2O(l)

This is the balanced redox reaction.