Q.Justify that the reaction: 2Cu2O(s) + Cu2S(s) → 6Cu(s) + SO2(g) is a redox reaction. Identify the species oxidised/reduced, which acts as an oxidant and which acts as a reductant.

Let’s start with something you already know from everyday life.

The intuition: what does “oxidation” really mean?

Think of a piece of iron left out in the rain. Over time, it turns into reddish-brown rust. Or think of a slice of apple turning brown when you leave it on the table. Or a fire burning wood to ash. In all these cases, something is combining with oxygen — that’s the original meaning of “oxidation.” The iron combines with oxygen from the air to form iron oxide (rust). The apple’s chemicals react with oxygen in the air. The wood burns because carbon in the wood combines with oxygen.

So the first, simplest idea: oxidation = adding oxygen. And the reverse — taking oxygen away — was called reduction. For example, if you heat iron oxide with carbon, the carbon steals the oxygen away, leaving pure iron. That’s reduction: removing oxygen.

But chemists soon realised this was too narrow. Many reactions that look like oxidation-reduction don’t involve oxygen at all. For instance, when sodium metal reacts with chlorine gas to make table salt, no oxygen is involved — yet the sodium clearly “rusts” in a sense, and the chlorine “steals” something from it.

So the definition had to be broadened.

The precise modern definition: electron transfer

Here’s the clean, exam-ready statement:

Oxidation is the loss of electrons by a substance.

Reduction is the gain of electrons by a substance.

They always happen together — you cannot have one without the other. That’s why we call them redox reactions (short for reduction-oxidation).

Let’s see this with the sodium-chlorine example:

• Sodium atom ($\text{Na}$) loses one electron to become $\text{Na}^+$. That’s oxidation.
• Chlorine atom ($\text{Cl}$) gains that electron to become $\text{Cl}^-$ . That’s reduction.

You can write the two halves separately:

Add them together:

That’s table salt.

How to spot a redox reaction without seeing electrons

You can’t watch electrons move directly. So chemists use oxidation numbers (also called oxidation states) — a bookkeeping system that tracks electrons.

Rules (simplified for first-time learners):

1. An atom in its elemental form has oxidation number 0.
2. A monatomic ion has oxidation number equal to its charge (e.g., $\text{Na}^+$ is +1, $\text{Cl}^-$ is -1).
3. Oxygen is usually -2 (except in peroxides).
4. Hydrogen is usually +1 (except in metal hydrides).
5. The sum of oxidation numbers in a neutral compound is 0; in a polyatomic ion, it equals the ion’s charge.

Then:

• Oxidation = increase in oxidation number.
• Reduction = decrease in oxidation number.

Example: Rusting of iron.

• Fe starts at 0 (elemental). In $\text{Fe}_2\text{O}_3$, each Fe is +3. So Fe’s oxidation number goes up from 0 to +3 → oxidation.
• O starts at 0 (in $\text{O}_2$). In $\text{Fe}_2\text{O}_3$, each O is -2. So O’s oxidation number goes down from 0 to -2 → reduction.

One more way to think about it …