Question

The reaction, \[\ce{\frac{1}{2} H2_{(g)} + AgCl_{(s)} -> H^+_{ (aq)} + Cl^-_{ (aq)} + Ag_{(s)}}\], occurs in the galvanic cell

Options

• \[\ce{Ag | AgCl_{(s)} | KCl (soln) || AgNO3 (soln) | Ag}\]

• \[\ce{Pt | H2_{(g)} | HCl (soln) || AgNO3 (soln) | Ag}\]

• \[\ce{Pt | H2_{(g)} | HCl (soln) || AgCl_{(s)} | Ag}\]

• \[\ce{Pt | H2_{(g)} | KCl (soln) || AgCl_{(s)} | Ag}\]

Answer 1

\[\ce{Pt | H2_{(g)} | HCl (soln) || AgCl_{(s)} | Ag}\]

Explanation:

The given reaction is a redox reaction.

\[\ce{\frac{1}{2} H2_{(g)} + AgCl_{(s)} -> H^+_{ (aq)} + Cl^-_{ (aq)} + Ag_{(s)}}\]

Where:

Anode (oxidation): \[\ce{H2 -> 2H+ + 2e-}\]

Cathode (reduction): \[\ce{AgCl + e- -> Ag + Cl-}\]

To represent this as a galvanic cell:

At the anode, hydrogen gas is oxidised.

\[\ce{Pt | H2_{(g)} | H+}\] (from HCl)

At the cathode, silver ion is reduced from AgCl.

\[\ce{AgCl_{(s)} | Ag}\]

So the correct cell representation is:

\[\ce{Pt | H2_{(g)} | HCl (soln) || AgCl_{(s)} | Ag}\]