Q.Balance the following redox reaction by ion - electron method:
MnO4^- (aq) + SO2

Redox Reaction Analysis – From Intuition to Precision

Imagine you have a piece of iron left out in the rain. Over time, it turns into reddish-brown rust. Something has clearly changed — the iron has combined with oxygen from the air. But what is actually happening at the atomic level? The iron atoms are losing something, and the oxygen atoms are gaining something. That "something" is electrons.

This exchange of electrons is the heart of every redox reaction. The word "redox" is a contraction of two processes that always happen together: reduction and oxidation.

The Core Intuition

Think of electrons as tiny, negatively charged particles that atoms can pass around like a hot potato. When an atom loses an electron, its charge becomes more positive. When an atom gains an electron, its charge becomes more negative.

• Oxidation = Loss of electrons (the atom becomes more positive)
• Reduction = Gain of electrons (the atom becomes more negative)

A simple mnemonic: OIL RIG — Oxidation Is Loss, Reduction Is Gain.

But here is the critical point: you cannot have one without the other. If one atom loses an electron, some other atom must gain it. That is why we call them redox reactions — they are a matched pair.

The Precise Statement

A redox reaction is any chemical reaction in which the oxidation states (or oxidation numbers) of atoms change. This change reflects the actual or formal transfer of electrons between species.

The two halves are:

1. Oxidation half-reaction: $X \rightarrow X^{n+} + n e^-$ (loss of $n$ electrons)
2. Reduction half-reaction: $Y + n e^- \rightarrow Y^{n-}$ (gain of $n$ electrons)

When combined, the electrons cancel out:

How to Analyse a Redox Reaction

You need a systematic method. Here is the step-by-step approach used in every Indian board exam:

Step 1: Assign oxidation numbers to every atom in the reaction. Rules (in order of priority):

• Free element = 0
• Monatomic ion = its charge
• Oxygen = -2 (except in peroxides where it is -1, and in OF$_2$ where it is +2)
• Hydrogen = +1 (except in metal hydrides where it is -1)
• Sum of oxidation numbers = charge of the species

Step 2: Identify which atoms change oxidation number. The atom that increases its oxidation number is oxidised (it lost electrons). The atom that decreases its oxidation number is reduced (it gained electrons).

Step 3: Write the two half-reactions — one for oxidation, one for reduction.

Step 4: Balance atoms and charge in each half-reaction (using H$_2$O, H$^+$, and OH$^-$ as needed for acidic/basic media).

Step 5: Multiply half-reactions so that electrons lost = electrons gained, then add them.

A Concrete Example

Consider the reaction between zinc metal and copper(II) sulphate:

Assign oxidation numbers:

• Zn (free element) = 0
• Cu in CuSO$_4$ = +2 (since SO$_4^{2-}$ has charge -2)
• Zn in ZnSO$_4$ = +2
• Cu (free element) = 0

What changed?

• Zn went from 0 to +2 → oxidised (lost 2 electrons)
• Cu went from +2 to 0 → reduced (gained 2 electrons)

Half-reactions:

• Oxidation: $\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$
• Reduction: $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$

The electrons cancel perfectly. The net reaction is simply the transfer of two electrons from Zn to Cu$^{2+}$.

Why This Matters …