Q.Can you store copper sulphate solutions in a zinc pot?

The key idea here is Galvanic Corrosion — when two dissimilar metals are in contact in the presence of an electrolyte, the more reactive metal corrodes.

1. Check the reactivity series: Zinc ($\ce{Zn}$) is more reactive than copper ($\ce{Cu}$).
2. In a zinc pot containing copper sulphate ($\ce{CuSO4}$) solution, zinc will displace copper from the solution:

3. This is a single displacement reaction. The zinc pot itself acts as the anode and corrodes, while copper metal deposits on the pot. …

The Core Idea: Galvanic Corrosion and Reactivity

This question is about galvanic corrosion — an electrochemical process that occurs when two dissimilar metals are in contact in the presence of an electrolyte. Here, the zinc pot is one metal, the copper sulphate solution is the electrolyte, and the copper ions in solution act as the second "metal" (they'll plate out as solid copper).

The driving force is the reactivity series of metals. Zinc is more reactive than copper — it sits higher in the series. This means zinc has a stronger tendency to lose electrons (oxidise) compared to copper. When you put them together in an electrolyte, zinc will spontaneously give up electrons, and copper ions will grab them.

This is a single displacement reaction — and it's the reason your zinc pot would be destroyed.

Step-by-Step Breakdown

1. Identify the reactants and their nature

You have solid zinc (the pot) and an aqueous solution of copper sulphate ($\text{CuSO}_4$). In solution, copper sulphate dissociates into $\text{Cu}^{2+}$ ions and $\text{SO}_4^{2-}$ ions. The zinc metal is a solid electrode.

2. Compare electrode potentials

The standard reduction potentials tell the story:

• $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$: $E^\circ = +0.34\ \text{V}$
• $\text{Zn}^{2+} + 2e^- \rightarrow \text{Zn}$: $E^\circ = -0.76\ \text{V}$

Zinc has a much more negative reduction potential — it's far more willing to lose electrons (oxidise) than copper is. In electrochemical terms, zinc acts as the anode (oxidation site) and copper ions act as the cathode (reduction site).

3. Write the half-reactions

Oxidation at the zinc surface:

Reduction of copper ions in solution:

4. Combine to see the net effect

The electrons lost by zinc are immediately taken up by copper ions. The net reaction:

This means:

• Zinc atoms leave the pot surface and go into solution as $\text{Zn}^{2+}$ ions — the pot corrodes.
• Copper ions come out of solution and deposit as solid copper metal — you'll see a reddish-brown coating on the pot. …

Method: Electrochemical Series / Standard Reduction Potential Analysis

This method uses the relative positions of metals in the electrochemical series to predict whether a redox reaction (and thus galvanic corrosion) will occur.

Steps

1. Identify the two metals involved

• Zinc (Zn) — the pot material
• Copper (Cu) — present as $\text{Cu}^{2+}$ ions in copper sulphate solution

2. Write the standard reduction potentials (from the electrochemical series)

• $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}(s)$ $E^\circ = +0.34\ \text{V}$
• $\text{Zn}^{2+} + 2e^- \rightarrow \text{Zn}(s)$ $E^\circ = -0.76\ \text{V}$

3. Determine which metal is more reactive (more easily oxidised)

• The more negative reduction potential means the metal is more easily oxidised.
• Zinc ($-0.76\ \text{V}$) is more reactive than copper ($+0.34\ \text{V}$).

4. Predict the spontaneous reaction

• Zinc will oxidise (lose electrons): $\text{Zn}(s) \rightarrow \text{Zn}^{2+} + 2e^-$
• Copper ions will reduce (gain electrons): $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}(s)$

Overall reaction:

$\text{Zn}(s) + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s)$ …

Common Mistakes: Storing Copper Sulphate in a Zinc Pot

✗ Mistake 1: Thinking "It's just a container — no reaction happens"

Why students make this: They treat the pot as an inert vessel, ignoring the chemical nature of the materials.

The truth: Zinc is more reactive than copper in the reactivity series. When $\text{CuSO}_4$ solution touches zinc, a single displacement reaction occurs:

How to avoid: Always check the reactivity series before storing a metal salt in a metal container. The container metal must be less reactive than the metal in the salt.

✗ Mistake 2: Confusing galvanic corrosion with simple rusting

Why students make this: They think corrosion only means iron rusting in air/water.

The truth: Galvanic corrosion happens when two dissimilar metals are in electrical contact in an electrolyte. Here:

• Zinc pot = anode (more reactive, corrodes)
• Copper ions in solution = cathode (less reactive, deposits as metal)
• CuSO₄ solution = electrolyte

The zinc sacrificially corrodes to deposit copper metal:

How to avoid: Remember: galvanic corrosion requires three things — anode, cathode, and electrolyte. Identify all three in any storage scenario.

✗ Mistake 3: Forgetting the pot gets destroyed

Why students make this: They focus only on the solution getting contaminated, not the container's fate.

The truth: The zinc pot dissolves over time. The reaction continues until either:

• All zinc is consumed (pot develops holes/leaks)
• All $\text{CuSO}_4$ is used up

How to avoid: Ask yourself: "Which material is being consumed?" The more reactive metal (zinc) is always the one that gets eaten away.

✗ Mistake 4: Thinking "Copper just settles at the bottom — no problem"

Why students make this: They see copper depositing and think it's harmless sediment.

The truth: The deposited copper accelerates further corrosion because:

• Copper metal on zinc creates a galvanic cell (Cu-Zn couple) …