Q.Assertion: Copper sulphate can be stored in zinc vessel.
Reason: Zinc is less reactive than copper.

The key idea here is Galvanic Corrosion — when two metals are in contact, the more reactive metal corrodes preferentially.

Reasoning:

1. In the reactivity series, zinc is more reactive than copper. Zinc can displace copper from its salt solution.
2. If copper sulphate is stored in a zinc vessel, the reaction $\ce{Zn + CuSO4 -> ZnSO4 + Cu}$ occurs. The vessel gets eaten away, and copper deposits on it. …

This question tests your understanding of the reactivity series of metals and displacement reactions — a core idea in electrochemistry and corrosion.

The key principle: a more reactive metal can displace a less reactive metal from its salt solution. If you store a salt solution in a container made of a more reactive metal, the container will react, dissolving itself and depositing the less reactive metal. That’s exactly what happens here.

Let’s break it down.

1. Check the assertion: “Copper sulphate can be stored in zinc vessel.” Zinc is above copper in the reactivity series. That means zinc is more reactive than copper. When zinc comes in contact with copper sulphate solution, a single displacement reaction occurs:

The zinc vessel would slowly dissolve, copper metal would deposit on it, and the solution would turn from blue to colourless (as $\text{Cu}^{2+}$ ions are replaced by $\text{Zn}^{2+}$). This is galvanic corrosion — the vessel gets eaten away.

So the assertion is false — you cannot store copper sulphate in a zinc vessel.

2. Check the reason: “Zinc is less reactive than copper.”

   This is the opposite of the truth. In the reactivity series, zinc is more reactive than copper. Zinc loses electrons more readily than copper does.

   So the reason is also false.

3. Evaluate the options:

   • (i) Both true, reason correct explanation — no, both false.
   • (ii) Both true, reason not correct explanation — no, both false. …

Method: Electrochemical Series / Standard Electrode Potential Comparison

This method uses the standard reduction potentials of the metals involved to predict whether a redox reaction will occur spontaneously.

Steps

1. Identify the two metals and their half-reactions

• Zinc (Zn) — the pot material

• Copper (Cu) — from copper sulphate solution

2. Determine which metal is more reactive (more negative $E^\circ$)

• Zinc has a more negative reduction potential → it is the stronger reducing agent (more easily oxidised).
• Copper has a more positive reduction potential → it is the weaker reducing agent (more easily reduced).

3. Predict the spontaneous reaction

• The more reactive metal (Zn) will oxidise (lose electrons).
• The less reactive metal ion ($\text{Cu}^{2+}$) will reduce (gain electrons).

Spontaneous reaction:

4. Calculate the cell potential to confirm …

Here’s a breakdown of the common mistakes students make on this galvanic corrosion / reactivity concept, and how to avoid each.

Mistake 1: Confusing the reactivity order of metals

• The error: Students think zinc is less reactive than copper, or they simply misremember the reactivity series.
• Why it happens: The reactivity series is often memorised without understanding the underlying trend (loss of electrons). Zinc loses electrons more readily than copper.
• How to avoid: Learn the reactivity series as a descending order of reactivity. A simple mnemonic: Potassium > Sodium > Calcium > Magnesium > Aluminium > Zinc > Iron > Tin > Lead > Hydrogen > Copper > Silver > Gold.
  • Zinc is above copper → more reactive.
• Key fact: A more reactive metal can displace a less reactive metal from its salt solution.

Mistake 2: Misinterpreting the “storage” condition

• The error: Students think “can be stored” means the container is inert, so they assume zinc is safe to use.
• Why it happens: They ignore the chemical reaction that would occur if the vessel is made of a more reactive metal.
• How to avoid: Always check: Is the vessel metal more reactive than the metal in the salt?
  • Zinc (vessel) is more reactive than copper (in copper sulphate).
  • Reaction: $\text{Zn} + \text{CuSO}_4 \rightarrow \text{ZnSO}_4 + \text{Cu}$
  • The vessel would corrode, copper would deposit, and the solution would be contaminated.
  • Conclusion: Copper sulphate cannot be stored in a zinc vessel.

Mistake 3: Thinking the reason is correct but not the explanation

• The error: Students pick option (ii) — “Both assertion and reason are true, but reason is not the correct explanation.”
• Why it happens: They know zinc is more reactive, but they think the reason statement (“Zinc is less reactive than copper”) is true because they’ve reversed the order.
• How to avoid: Verify the reason statement independently.
  • The reason says: “Zinc is less reactive than copper.”
  • This is false (zinc is more reactive).
  • So the reason is false → options (i) and (ii) are eliminated.

Mistake 4: Forgetting the concept of galvanic corrosion

• The error: Students treat this as a simple “can it be stored?” question without linking it to corrosion.
• Why it happens: They don’t connect the displacement reaction to the formation of a galvanic cell (zinc anode, copper cathode) in the presence of an electrolyte.
• How to avoid: Understand the mechanism:
  • When zinc is in contact with copper sulphate solution, a galvanic cell forms.
  • Zinc (more reactive) acts as the anode and corrodes (loses electrons).
  • Copper ions gain electrons and deposit on the zinc surface.
  • This is galvanic corrosion — the vessel gets eaten away.

Mistake 5: Choosing option (iii) — “Assertion is true, reason is false” …