The reaction

$$\frac{1}{2} \mathrm{H}_{2}(\mathrm{~g})+\mathrm{AgCl}(\mathrm{s}) \rightleftharpoons \mathrm{H}^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s})$$

occurs in which of the given galvanic cell.

<p>The provided reaction is:</p> <p>$$\frac{1}{2} \mathrm{H}_{2}(\mathrm{~g})+\mathrm{AgCl}(\mathrm{s}) \rightleftharpoons \mathrm{H}^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s})$$</p> <p>This reaction involves the following half-reactions:</p> <ol> <li>Oxidation of hydrogen gas to H+ ions: $$\frac{1}{2} \mathrm{H}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{H}^{+}(\mathrm{aq}) + e^-$$</li><br/> <li>Reduction of AgCl to Ag: $$\mathrm{AgCl}(\mathrm{s}) + e^- \rightleftharpoons \mathrm{Ag}(\mathrm{s}) + \mathrm{Cl}^{-}(\mathrm{aq})$$</li> </ol> <p>Looking at the options provided:</p> <p>Option A: Doesn&#39;t involve H₂ gas, so it can&#39;t be correct.</p> <p>Option B: This includes the necessary elements - H₂, AgCl, and Ag. </p> <p>Option C: Doesn&#39;t involve AgCl, so it can&#39;t be correct.</p> <p>Option D: Also includes the necessary elements - H₂, AgCl, and Ag.</p> <p>However, looking closely, we can see that Option B represents the galvanic cell for this reaction. The reaction requires the oxidation of H₂ to H⁺, which occurs at the anode. The reaction also requires the reduction of AgCl to Ag and Cl-, which occurs at the cathode.</p> <p>In Option B, the anode (on the left) is where H₂ is being oxidized to H⁺. The cathode (on the right) is where AgCl is reduced to Ag and Cl^-. The salt bridge or ion exchange component is HCl, which allows for the flow of ions to balance charge in the cell.</p> <p>Therefore, the reaction occurs in the galvanic cell represented by Option B.</p>