(i) Cu (s) does not dissolve in HCl. Why? (ii) Can you keep copper sulphate solution in a vessel of zinc? Explain with reason.

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> [!TLDR]
> Copper's positive electrode potential means it cannot displace $H_2$ from HCl; zinc's very negative potential means it displaces copper from $CuSO_4$, so copper sulphate cannot be stored in a zinc vessel.
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> **(i) Why Cu(s) does not dissolve in HCl.** Dissolving a metal in a non-oxidising acid like HCl requires the metal to reduce $H^+$ to $H_2$, i.e. the metal must be *above* hydrogen in the activity series (negative $E^{\circ}$).
> $$E^{\circ}_{Cu^{2+}/Cu}=+0.34\ 	ext{V}\;(	ext{positive})$$
> Since copper's standard reduction potential is positive (it lies *below* hydrogen), $H^+$ cannot oxidise copper. The reaction $Cu + 2H^+ 	o Cu^{2+} + H_2$ has a negative cell potential and does not occur, so **Cu does not dissolve in HCl** (it dissolves only in oxidising acids such as $HNO_3$).
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> **(ii) Copper sulphate solution in a zinc vessel.** Compare the metals:
> $$E^{\circ}_{Zn^{2+}/Zn}=-0.76\ 	ext{V},\qquad E^{\circ}_{Cu^{2+}/Cu}=+0.34\ 	ext{V}$$
> Zinc is more easily oxidised (more negative potential), so it will **displace copper** from the solution:
> $$Zn(s) + CuSO_4(aq) \longrightarrow ZnSO_4(aq) + Cu(s)\qquad E^{\circ}_{cell}=+1.10\ 	ext{V}$$ …

Element	Half-Reaction (Reduction)	$E^\ominus$ (V)
Lithium	$\text{Li}^+ + e^- \rightarrow \text{Li}$	$-3.04$
Potassium	$\text{K}^+ + e^- \rightarrow \text{K}$	$-2.93$
Calcium	$\text{Ca}^{2+} + 2e^- \rightarrow \text{Ca}$	$-2.87$
Sodium	$\text{Na}^+ + e^- \rightarrow \text{Na}$	$-2.71$
Magnesium	$\text{Mg}^{2+} + 2e^- \rightarrow \text{Mg}$	$-2.37$
Aluminium	$\text{Al}^{3+} + 3e^- \rightarrow \text{Al}$	$-1.66$
Zinc	$\text{Zn}^{2+} + 2e^- \rightarrow \text{Zn}$	$-0.76$
Iron	$\text{Fe}^{2+} + 2e^- \rightarrow \text{Fe}$	$-0.44$
Tin	$\text{Sn}^{2+} + 2e^- \rightarrow \text{Sn}$	$-0.14$
Lead	$\text{Pb}^{2+} + 2e^- \rightarrow \text{Pb}$	$-0.13$
Hydrogen	$\text{2H}^+ + 2e^- \rightarrow \text{H}_2$	$0.00$
Copper	$\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$	$+0.34$
Silver	$\text{Ag}^+ + e^- \rightarrow \text{Ag}$	$+0.80$
Gold	$\text{Au}^{3+} + 3e^- \rightarrow \text{Au}$	$+1.50$

Q.(i) Cu (s) does not dissolve in HCl. Why?

Concept understanding — Electrochemical Series

The Intuition: A "Tug-of-War" for Electrons

The Precise Statement

What the Series Tells You

1. Predicting Reactivity

2. Displacement Reactions

3. Direction of Redox Reactions …

Q.(i) Cu (s) does not dissolve in HCl. Why?

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