Assertion: Copper sulphate can be stored in zinc vessel. Reason: Zinc is less reactive than copper. (i) Both a

NCERT Exemplar · Q68

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

(i) Both assertion and reason are true and the reason is the correct explanation of assertion.

(ii) Both assertion and reason are true and the reason is not the correct explanation of assertion.

(iii) Assertion is true but the reason is false.

(iv) Both assertion and reason are false.

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Concept understanding — Galvanic Corrosion

Galvanic Corrosion: From Intuition to Precision

Imagine you have two different metals — say, a copper pipe and an iron nail — and you connect them with a wire, then dip both into a bucket of salt water. If you come back a few hours later, the iron nail will be badly rusted, while the copper pipe will look almost untouched. Why?

The answer is galvanic corrosion. It is the accelerated corrosion of one metal when it is in electrical contact with a different metal in the presence of an electrolyte (like water with dissolved salts).

The Intuition: A "Battery" That Eats Metal

Think of a simple battery: you have two different metals (electrodes) and a chemical solution (electrolyte). One metal wants to give away electrons (it gets eaten away), and the other wants to accept them (it stays protected). That is exactly what happens in galvanic corrosion.

The more reactive metal (the one that "wants" to corrode) becomes the anode. It loses electrons and dissolves into the electrolyte — that is the corrosion you see.
The less reactive metal becomes the cathode. It does not corrode; instead, it accepts electrons from the anode, often causing the electrolyte near it to become alkaline or to produce hydrogen gas.

The key point: the two metals do not need to be physically touching. They just need electrical contact (through a wire or direct contact) and a continuous electrolyte (water, soil, concrete, etc.) to complete the circuit.

The Precise Statement

Important

Galvanic corrosion is the electrochemical process in which a more active metal (the anode) corrodes preferentially when electrically coupled to a less active metal (the cathode) in the presence of an electrolyte. The driving force is the difference in their electrode potentials.

The Galvanic Series: The "Who Eats Whom" Chart

Not all metal pairs corrode equally. The galvanic series ranks metals and alloys by their tendency to corrode in seawater (a common electrolyte). The more negative (active) a metal is, the more likely it is to be the anode and corrode.

Here is a simplified version of the series (from most active/anodic to most noble/cathodic):

Metal / Alloy	Relative Activity
Magnesium	Most active (anodic)
Zinc
Aluminium
Cadmium
Mild steel / Iron
Stainless steel (active)
Tin
Lead
Copper
Nickel
Stainless steel (passive)
Silver
Titanium
Gold / Platinum	Most noble (cathodic)

Watch out

A common mistake: students think the larger metal always corrodes. In reality, it is the more active metal that corrodes, regardless of size. However, the area ratio matters enormously — a small anode coupled to a large cathode corrodes very fast (like a tiny iron rivet holding a huge copper plate).

The Three Conditions for Galvanic Corrosion

For galvanic corrosion to occur, all three must be present:

Two dissimilar metals (or the same metal in different environments, e.g., a steel pipe in soil vs. in air).
Electrical contact between them (direct physical contact or through a wire).
An electrolyte bridging them (water, moisture, soil, concrete, etc.).

Remove any one, and galvanic corrosion stops.

Real-World Examples

The Statue of Liberty: The copper skin was originally separated from the iron framework by asbestos cloth. When the cloth degraded, the iron (anode) corroded rapidly because it was coupled to the huge copper (cathode) surface. …

Why this formula?

Galvanic Corrosion: Why the Key Formulas Hold

Galvanic corrosion occurs when two dissimilar metals are electrically connected in the presence of an electrolyte. The key formula that governs this is the mixed potential theory, which leads to the galvanic current and corrosion rate expressions.

Let's build the reasoning step-by-step.

1. The Core Idea: Two Electrodes, One Circuit

When metals M₁ (more active, e.g., zinc) and M₂ (more noble, e.g., copper) are connected:

M₁ acts as the anode — it oxidizes (corrodes):

\text{M}_1 \rightarrow \text{M}_1^{n+} + n e^-

M₂ acts as the cathode — it reduces something (e.g., oxygen or H⁺):

\text{O}_2 + 2\text{H}_2\text{O} + 4e^- \rightarrow 4\text{OH}^- \quad \text{(in neutral/alkaline)}

2\text{H}^+ + 2e^- \rightarrow \text{H}_2 \quad \text{(in acidic)}

The two metals are electrically connected (via a wire or direct contact), and the electrolyte completes the circuit. Electrons flow from M₁ to M₂.

2. The Mixed Potential: Why It Exists

Each metal, when alone in the electrolyte, has its own open-circuit potential (OCP) — the equilibrium potential for its half-reaction. For M₁, it's $E_{\text{corr},1}$ ; for M₂, it's $E_{\text{corr},2}$ .

When connected, the system cannot stay at two different potentials. The entire metal couple must reach a single potential — the mixed potential $E_{\text{mix}}$ .

$E_{\text{mix}}$ lies between $E_{\text{corr},1}$ and $E_{\text{corr},2}$ .
At $E_{\text{mix}}$ , the total anodic current from M₁ equals the total cathodic current from M₂ (charge conservation):

I_{\text{anode}} = I_{\text{cathode}}

This is the fundamental equation of galvanic corrosion.

3. Deriving the Galvanic Current

Assume each electrode follows Butler-Volmer kinetics (for activation-controlled reactions). For the anode (M₁), the anodic current density $i_a$ at potential $E$ is:

i_a = i_{0,1} \exp\left( \frac{\alpha_a F (E - E_{0,1})}{RT} \right)

For the cathode (M₂), the cathodic current density $i_c$ is:

i_c = i_{0,2} \exp\left( -\frac{\alpha_c F (E - E_{0,2})}{RT} \right)

Where:

$i_{0,1}, i_{0,2}$ = exchange current densities
$\alpha_a, \alpha_c$ = transfer coefficients (typically ~0.5)
$F$ = Faraday constant
$R$ = gas constant
$T$ = temperature
$E_{0,1}, E_{0,2}$ = standard reduction potentials

At the mixed potential $E_{\text{mix}}$ :

I_{\text{galvanic}} = A_1 \cdot i_a(E_{\text{mix}}) = A_2 \cdot i_c(E_{\text{mix}})

Where $A_1$ and $A_2$ are the surface areas of the anode and cathode.

Why this holds: The net current from the anode must exactly balance the net current consumed at the cathode — otherwise, charge would accumulate, which is impossible in a steady-state circuit.

4. The Corrosion Rate Formula

The corrosion rate (mass loss per time) of the anode is given by Faraday's law:

\text{Corrosion rate} = \frac{I_{\text{galvanic}} \cdot M}{n \cdot F \cdot \rho}

Where:

$M$ = molar mass of the anode metal
$n$ = number of electrons transferred per atom
$\rho$ = density of the metal
$F$ = Faraday constant (96,485 C/mol)

Why this holds: Each mole of metal oxidized releases $n$ moles of electrons. The total charge passed $Q = I_{\text{galvanic}} \cdot t$ corresponds to moles of metal lost:

\text{moles lost} = \frac{Q}{nF} = \frac{I_{\text{galvanic}} \cdot t}{nF}

Multiply by $M/\rho$ to get volume or thickness loss.

5. The Area Effect: Why It Matters

From the mixed potential equation:

A_1 \cdot i_a(E_{\text{mix}}) = A_2 \cdot i_c(E_{\text{mix}})

If the cathode area $A_2$ is large relative to the anode area $A_1$ , then $i_a(E_{\text{mix}})$ must be large to balance the current. This means:

Small anode + large cathode → severe galvanic corrosion (high current density on the anode). …

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The assertion is false because zinc is more reactive than copper, so zinc would displace copper from copper sulphate, corroding the vessel. The reason is also false (zinc is more reactive, not less). Hence both are false — option (iv).

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.

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:

\text{Zn (s)} + \text{CuSO}_4 \text{(aq)} \rightarrow \text{ZnSO}_4 \text{(aq)} + \text{Cu (s)}

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.

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.
Evaluate the options:
- (i) Both true, reason correct explanation — no, both false.
- (ii) Both true, reason not correct explanation — no, both false. …

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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” …

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Same / Similar Concept — real previous-year questions on the same or a closely similar concept, not this exact question.

KCET 2025Set D-41 markMCQ
$\mathrm{Cmol}^{-1}$
›Reveal solutionSolution
The charge needed per mole of metal is just $n \times F$ , where $n$ is the cation's charge and $F = 96500\,\mathrm{C\,mol^{-1}}$ — so simply multiply $96500$ by $1, 2, 3, 4$ .

Step 1 — The governing law.

The cathodic reduction is
$\mathrm{M}^{n+} + n e^- \rightarrow \mathrm{M}$
One mole of $\mathrm{M}^{n+}$ therefore consumes exactly $n$ moles of electrons. Faraday's constant is the charge on one mole of electrons,
$F = 96500\ \mathrm{C\,mol^{-1}}$
so the quantity of electricity required per mole of metal deposited is
$Q = nF$
Step 2 — Compute for each cation.

(a) $\mathrm{Ag^+}$ , $n = 1$ : $Q = 1 \times 96500 = 96500\ \mathrm{C\,mol^{-1}}$ → (iii)

(b) $\mathrm{Mg^{2+}}$ , $n = 2$ : $Q = 2 \times 96500 = 193000\ \mathrm{C\,mol^{-1}}$ → (iv)

(c) $\mathrm{Al^{3+}}$ , $n = 3$ : $Q = 3 \times 96500 = 289500\ \mathrm{C\,mol^{-1}}$ → (ii)

(d) $\mathrm{Ti^{4+}}$ , $n = 4$ : $Q = 4 \times 96500 = 386000\ \mathrm{C\,mol^{-1}}$ → (i)

Step 3 — Assemble the matching. …
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KCET 2024Set B-21 markMCQ

\approx 5\%

›Reveal solution

Recall the composition of Misch metal: ~95% lanthanoid + ~5% iron, with traces of S, C, Ca and Al.

Step 1 — What Misch metal is.

Misch metal is a commercially important lanthanoid alloy. Its composition is:

~95% lanthanoid metal (mostly cerium, with lanthanum, neodymium and praseodymium),
~5% iron,
traces of S, C, Ca and Al.

Step 2 — Identify the transition element.

The question asks specifically for the transition element present at the ~5% level. Screening the options against the $d$ -block:

Option	Element	Block
(A)	Mg	$s$ -block (alkaline earth) — not a transition metal
(B)	Fe	$3d$ transition series ✓
(C)	Zn	$3d$ series, but $d^{10}$ in all its compounds — and not the 5% component
(D)	Co	$3d$ series, but not a component of Misch metal

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COMEDK 2024Set 2024-A1 markMCQ
$\mathrm{E}^{\circ}$
›Reveal solutionSolution
To prevent iron corrosion by coating, the coating metal must be more easily oxidized (more negative reduction potential) than iron, so it acts as a sacrificial anode. Only element A ( $E^\circ = -2.37\ \mathrm{V}$ ) satisfies this, making option (D) correct.

Concept & Intuition

Corrosion of iron (rusting) is an electrochemical process where iron is oxidized:
$\mathrm{Fe} \rightarrow \mathrm{Fe}^{2+} + 2e^- \quad (E^\circ = -0.44\ \mathrm{V})$
If we coat iron with another metal, we want that metal to oxidize instead of iron — it “sacrifices” itself. This happens if the coating metal has a more negative reduction potential than iron, meaning it is a stronger reducing agent (more easily oxidized). A metal with a less negative or positive reduction potential would actually cause iron to corrode faster (it would act as a cathode, forcing iron to be the anode).

Step-by-step reasoning

Identify iron’s tendency to oxidize

Iron’s reduction potential is $E^\circ_{\mathrm{Fe}^{2+}/\mathrm{Fe}} = -0.44\ \mathrm{V}$ . The more negative this value, the easier it is for the metal to lose electrons (oxidize). So iron itself is moderately prone to oxidation.

Compare each candidate metal’s reduction potential

Element A: $E^\circ_{\mathrm{A}^{2+}/\mathrm{A}} = -2.37\ \mathrm{V}$ — much more negative than iron.

Element B: $E^\circ_{\mathrm{B}^{2+}/\mathrm{B}} = -0.15\ \mathrm{V}$ — less negative than iron.

Element C: $E^\circ_{\mathrm{C}^{2+}/\mathrm{C}} = +0.34\ \mathrm{V}$ — positive, meaning it is very hard to oxidize.

Determine which metal will oxidize preferentially

For a coating to protect iron, the coating metal must have a more negative reduction potential than iron. That way, when an electrochemical cell forms (e.g., with moisture and oxygen), the coating metal becomes the anode and corrodes, while iron remains the cathode and is protected.

A ( $-2.37\ \mathrm{V}$ ) is more negative than Fe ( $-0.44\ \mathrm{V}$ ) → A will protect iron.

B ( $-0.15\ \mathrm{V}$ ) is less negative than Fe → B is harder to oxidize than iron, so iron would corrode first.

C ( $+0.34\ \mathrm{V}$ ) is positive → C is very noble; it would actually accelerate iron corrosion (iron becomes the sacrificial anode). …
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COMEDK 2024Set 2024-E1 markMCQ
$\mathrm{O}_2(\mathrm{~g})+2 \mathrm{H}_2 \mathrm{O}(\mathrm{l})+4 \mathrm{e} \rightarrow 4 \mathrm{OH}^{-}$
›Reveal solutionSolution
The cathodic (reduction) reaction in rusting is $O_2+4H^+ +4e^-\rightarrow 2H_2O$ with $E^0=+1.23\,V$ — statement (B).

Electrochemical mechanism of rusting (NCERT):

Anode (oxidation): $Fe(s)\rightarrow Fe^{2+}+2e^-$ , $E^0_{(Fe^{2+}/Fe)}=-0.44\,V$ .

Cathode (reduction): in the presence of $H^+$ (from dissolved $CO_2$ /water) and atmospheric oxygen, $O_2(g)+4H^+(aq)+4e^-\rightarrow 2H_2O(l)$ , $E^0=+1.23\,V$ .

Overall: $2Fe+O_2+4H^+\rightarrow 2Fe^{2+}+2H_2O$ ; the $Fe^{2+}$ is then further oxidised to hydrated $Fe_2O_3$ (rust). …
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Q.Assertion: Copper sulphate can be stored in zinc vessel.
Reason: Zinc is less reactive than copper.

Concept understanding — Galvanic Corrosion

Galvanic Corrosion: From Intuition to Precision

The Intuition: A "Battery" That Eats Metal

The Precise Statement

The Galvanic Series: The "Who Eats Whom" Chart

The Three Conditions for Galvanic Corrosion

Real-World Examples

Why this formula?

Galvanic Corrosion: Why the Key Formulas Hold

1. The Core Idea: Two Electrodes, One Circuit

2. The Mixed Potential: Why It Exists

3. Deriving the Galvanic Current

4. The Corrosion Rate Formula

5. The Area Effect: Why It Matters

Method: Electrochemical Series / Standard Electrode Potential Comparison

Steps

1. Identify the two metals and their half-reactions

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

3. Predict the spontaneous reaction

4. Calculate the cell potential to confirm …

Mistake 1: Confusing the reactivity order of metals

Mistake 2: Misinterpreting the “storage” condition

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

Mistake 4: Forgetting the concept of galvanic corrosion

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

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

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

Unlock everything free for 14 days