The overall cell reaction is: — Electrochemistry Chemistry Question

💡 Solution & Explanation

Step 1 - Identify the Electrodes from the Cell Notation The electrochemical cell is represented using standard IUPAC notation as: $$\text{Pt}(s) \mid \ce{H2}(g) \mid \ce{HCl}(aq) \mid \ce{AgCl}(s) \mid \ce{Ag}(s)$$ According to electrochemical cell conventions: * The left-hand side represents the **anode** (where oxidation occurs): $$\text{Pt}(s) \mid \ce{H2}(g) \mid \ce{H^+}(aq)$$ * The right-hand side represents the **cathode** (where reduction occurs): $$\ce{Cl^-(aq) \mid AgCl(s) \mid Ag(s)}$$ Both electrodes are in contact with the aqueous hydrochloric acid ($\ce{HCl}$) electrolyte, which provides both $\ce{H^+}$ and $\ce{Cl^-}$ ions. Step 2 - Write the Anode Half-Reaction (Oxidation) At the anode, hydrogen gas is bubbled over the platinum surface and undergoes oxidation to form hydrogen ions, releasing electrons into the external circuit: $$\ce{H2(g) -> 2H^+(aq) + 2e^-}$$ Step 3 - Write the Cathode Half-Reaction (Reduction) The cathode is a silver-silver chloride electrode (a metal-insoluble salt-anion electrode). During reduction, solid silver chloride accepts electrons at the silver metal surface to form metallic silver and release free chloride ions into the solution: $$\ce{AgCl(s) + e^- -> Ag(s) + Cl^-(aq)}$$ To balance the electrons transferred in the anode reaction (which is $2\text{e}^-$), we multiply the cathode half-reaction by $2$: $$\ce{2AgCl(s) + 2e^- -> 2Ag(s) + 2Cl^-(aq)}$$ Step 4 - Determine the Overall Cell Reaction To find the net cell reaction, we add the oxidation half-reaction and the balanced reduction half-reaction together: $$\text{Anode: } \ce{H2(g) -> 2H^+(aq) + 2e^-}$$ $$\text{Cathode: } \ce{2AgCl(s) + 2e^- -> 2Ag(s) + 2Cl^-(aq)}$$ $$\text{Overall Reaction: } \ce{H2(g) + 2AgCl(s) -> 2H^+(aq) + 2Cl^-(aq) + 2Ag(s)}$$ Since hydrochloric acid ($\ce{HCl}$) is a strong electrolyte that exists completely as dissociated ions ($\ce{H^+}$ and $\ce{Cl^-}$) in aqueous solution, we can combine the ionic species on the product side to represent the molecular reaction: $$\ce{H2(g) + 2AgCl(s) -> 2HCl(aq) + 2Ag(s)}$$ Step 5 - Evaluate and Explain Each Option * **Option (A) is correct:** As derived in Step 4, this equation correctly represents the stoichiometry and species involved in the overall chemical change in the cell. * **Option (B) is incorrect:** This equation represents a cell where free silver ions ($\ce{Ag^+}$) are reduced to silver metal. However, in this cell, the concentration of free $\ce{Ag^+}$ is extremely low and the primary reacting species is the solid salt $\ce{AgCl(s)}$. * **Option (C) is incorrect:** This is a simple precipitation reaction, not a redox process that can generate a standard cell potential. * **Option (D) is incorrect:** This is the direct gas-phase reaction of hydrogen and chlorine, which does not represent the chemistry of the silver-silver chloride electrode. $$\text{Correct Option: } \boxed{A}$$