Ammonia is always added in this reaction. Which of the following must be incorrect? — Electrochemistry Chemistry Question

💡 Solution & Explanation

Step 1 - Understand the Chemical Principles of Tollen's Reagent (Option A) Tollen's reagent is an aqueous solution of silver nitrate ($\ce{AgNO3}$) mixed with ammonium hydroxide ($\ce{NH4OH}$), which acts as a source of ammonia ($\ce{NH3}$). Free silver ions ($\ce{Ag^+}$) are highly unstable in basic media because they react with hydroxide ions ($\ce{OH^-}$) to precipitate brown silver oxide ($\ce{Ag2O}$): $$\ce{2Ag^+(aq) + 2OH^-(aq) -> Ag2O(s) + H2O(l)}$$ To prevent this precipitation, ammonia is added. It acts as a ligand and coordinates with silver ions to form a stable, soluble coordination complex known as the diamminesilver(I) complex: $$\ce{Ag^+(aq) + 2NH3(aq) <=> [Ag(NH3)2]^+(aq)}$$ Since ammonia combines with $\ce{Ag^+}$ to form a complex, **Option (A) represents a correct chemical statement** and is therefore not the incorrect choice. Step 2 - Analyze the Oxidizing Power of the Species (Option B) The oxidizing strength of a chemical species is determined by its tendency to undergo reduction, which is measured quantitatively by its standard reduction potential ($E^\circ_{\text{red}}$). A higher, more positive standard reduction potential indicates a stronger tendency to accept electrons, making the species a stronger oxidizing agent. We are given the following standard reduction potentials at $298\text{ K}$: * For free silver ions: $$\ce{Ag^+(aq) + e^- -> Ag(s)} \quad E^\circ_{\text{red}} = 0.8\text{ V}$$ * For the diamminesilver(I) complex: $$\ce{Ag(NH3)2^+(aq) + e^- -> Ag(s) + 2NH3(aq)} \quad E^\circ_{\text{red}} = 0.337\text{ V}$$ Comparing the two standard reduction potential values: $$E^\circ_{\text{red}}(\ce{Ag^+}) = 0.8\text{ V} > E^\circ_{\text{red}}(\ce{Ag(NH3)2^+}) = 0.337\text{ V}$$ Because the standard reduction potential of free $\ce{Ag^+}$ is significantly greater than that of the complex $\ce{Ag(NH3)2^+}$, free $\ce{Ag^+}$ is a much stronger oxidizing agent than the complex. The complexation of silver ions with ammonia stabilizes the $+1$ oxidation state of silver in solution, which significantly lowers its tendency to be reduced, thereby reducing its oxidizing power. Therefore, the statement in **Option (B)** "$\ce{Ag(NH3)2^+}$ is a stronger oxidizing agent than $\ce{Ag^+}$" is **thermodynamically incorrect (false)**. Since we are looking for the incorrect statement, this is our correct option. Step 3 - Analyze the Formation of the Silver Salt of Gluconic Acid (Option C) During the redox reaction, glucose ($\ce{C6H12O6}$) is oxidized to gluconic acid ($\ce{C6H12O7}$). Since the reaction takes place in an alkaline medium provided by ammonium hydroxide/ammonia, the carboxylic acid group of gluconic acid is neutralized by the base to form gluconate ions: $$\ce{C6H12O6 + 2[Ag(NH3)2]^+ + 3OH^- -> C6H11O7^- + 2Ag(s) + 4NH3 + 2H2O}$$ In the presence of silver species, the silver salt of gluconic acid (silver gluconate) can be formed or exists in equilibrium. Therefore, **Option (C) represents a correct chemical statement**. Step 4 - Analyze the Effect of Ammonia on the Glucose Electrode Potential (Option D) The standard half-cell oxidation of glucose to gluconic acid is represented by: $$\ce{C6H12O6 + H2O -> C6H12O7 + 2H^+ + 2e^-}$$ According to the Nernst equation, the electrode potential of this hydrogen-ion-dependent half-cell is extremely sensitive to pH: $$E_{\text{ox}} = E^\circ_{\text{ox}} - \frac{0.0592}{2}\log_{10}\left(\frac{[\ce{C6H12O7}][\ce{H^+}]^2}{[\ce{C6H12O6}]}\right)$$ The addition of ammonia ($\ce{NH3}$), which is a weak base, neutralizes hydronium ions ($\ce{H^+}$) to form ammonium ions ($\ce{NH4^+}$), dramatically raising the pH of the medium. This reduction in $[\ce{H^+}]$ shifts the equilibrium of the oxidation reaction to the right, significantly increasing the oxidation potential ($E_{\text{ox}}$) and making the overall reaction thermodynamically more spontaneous. Because ammonia modifies the concentration of the potential-determining ions ($\ce{H^+}$) of this redox couple, it directly determines the non-standard electrode potential. In biochemical and organic systems, standard potentials are often defined at a specific pH (such as pH 7 or alkaline pH), which is directly controlled by the basic buffer system. Thus, **Option (D) represents a correct statement**. Step 5 - Conclusion Only statement (B) is thermodynamically and chemically incorrect. $$\text{Correct Option: } \boxed{B}$$