Write the electrolytic products of electrolysis of (i) aq \(\mathrm{NaOH}_{\text {(conc) }}\) (Pt electrodes) (ii) aq \(\mathrm{NaCl}_{(\text {conc })}(\) Pt electrodes \()\) (iii) aq \(\mathrm{AgNO}_{3(c o n c)}\) (Ag electrodes)

To determine the electrolytic products for this solution, we first identify the potential anodic and cathodic reactions: Possible anodic reactions: 1. Oxidation of hydroxide ions: \(\mathrm{4OH}^{-} \rightarrow \mathrm{O}_{2} + \mathrm{2H}_{2}\mathrm{O} + \mathrm{4e}^{-}\) 2. Oxidation of water: \(\mathrm{2H}_{2}\mathrm{O} \rightarrow \mathrm{O}_{2} + \mathrm{4H}^{+} + \mathrm{4e}^{-}\) Possible cathodic reactions: 1. Reduction of water: \(\mathrm{2H}_{2}\mathrm{O} + \mathrm{2e}^{-} \rightarrow \mathrm{H}_{2} + \mathrm{2OH}^{-}\) Comparing the standard electrode potentials, it is more likely for hydroxide ions to be oxidized at the anode, and for water to be reduced at the cathode. Thus, the electrolytic products are: Anode: Oxygen gas (\(\mathrm{O}_{2}\)) Cathode: Hydrogen gas (\(\mathrm{H}_{2}\))

Similarly, we identify the possible anodic and cathodic reactions for this solution: Possible anodic reactions: 1. Oxidation of chloride ions: \(\mathrm{2Cl}^{-} \rightarrow \mathrm{Cl}_{2} + \mathrm{2e}^{-}\) 2. Oxidation of water: \(\mathrm{2H}_{2}\mathrm{O} \rightarrow \mathrm{O}_{2} + \mathrm{4H}^{+} + \mathrm{4e}^{-}\) Possible cathodic reactions: 1. Reduction of water: \(\mathrm{2H}_{2}\mathrm{O} + \mathrm{2e}^{-} \rightarrow \mathrm{H}_{2} + \mathrm{2OH}^{-}\) 2. Reduction of sodium ions: \(\mathrm{Na}^{+} + \mathrm{e}^{-} \rightarrow \mathrm{Na}\) Since chloride ions have a lower standard electrode potential than water, they will be preferentially oxidized at the anode. At the cathode, water reduction is more likely due to the high standard electrode potential required for sodium reduction in aqueous solution. Thus, the electrolytic products are: Anode: Chlorine gas (\(\mathrm{Cl}_{2}\)) Cathode: Hydrogen gas (\(\mathrm{H}_{2}\))

Finally, we identify the possible anodic and cathodic reactions for this solution: Possible anodic reactions: 1. Oxidation of silver electrode: \(\mathrm{Ag} \rightarrow \mathrm{Ag}^{+} + \mathrm{e}^{-}\) Possible cathodic reactions: 1. Reduction of silver ions: \(\mathrm{Ag}^{+} + \mathrm{e}^{-} \rightarrow \mathrm{Ag}\) 2. Reduction of water: \(\mathrm{2H}_{2}\mathrm{O} + \mathrm{2e}^{-} \rightarrow \mathrm{H}_{2} + \mathrm{2OH}^{-}\) 3. Reduction of nitrate ions: \(\mathrm{NO}_{3}^{-} + \mathrm{2H}_{2}\mathrm{O} + \mathrm{2e}^{-} \rightarrow \mathrm{NO}_{2}^{-} + \mathrm{4OH}^{-}\) In this case, the standard electrode potential for the reduction of silver ions is very high, making it the most likely reaction to occur at the cathode. The anodic reaction will be the oxidation of the silver electrode. Thus, the electrolytic products are: Anode: None (silver ions enter the solution, but no product forms) Cathode: Silver metal (\(\mathrm{Ag}\)) is deposited on the electrode