The standard reduction potential of Ag^+/Ag electrode is +0.80 V. Cho…

The correct answer is Option (2) → \(Ag^+/Ag\) couple is a weaker reducing agent than \(H^+/H_2\) couple

Standard electrode potentials are measured against the standard hydrogen electrode, whose E° is taken as 0.00 V. A positive E° means that the ion is reduced more easily than \( \ce{H^+} \); the metal is then less easily oxidised than hydrogen, so the couple is a weaker reducing agent than the \( \ce{H^+/H2} \) couple. A negative E° means the opposite.

For silver, E° = +0.80 V. So \( \ce{Ag^+} \) is a better oxidising agent than \( \ce{H^+} \), and Ag is a weaker reducing agent than \( \ce{H2} \). In a cell with the SHE, silver is the cathode:

    \( \displaystyle \text{Anode:}\quad \ce{H2 -> 2H^+ + 2e^-} \)

    \( \displaystyle \text{Cathode:}\quad \ce{2Ag^+ + 2e^- -> 2Ag} \)

    \( \displaystyle \text{Cell:}\quad \ce{H2 + 2Ag^+ -> 2H^+ + 2Ag}, \qquad E^\circ_{\text{cell}} = 0.80 - 0.00 = +0.80\ \mathrm{V} \)

So the correct statement is that the \( \ce{Ag^+/Ag} \) couple is a weaker reducing agent than the \( \ce{H^+/H2} \) couple.

Why the other options are incorrect

• Option (1) \(Ag^+/Ag\) couple is a stronger reducing agent than \(H^+/H_2\) couple — Incorrect. A stronger reducing agent than \( \ce{H2} \) would need a negative E°, as for Zn (−0.76 V).

• Option (3) \(Ag\) can be easily oxidised as compared to \(H_2\) — Incorrect. The positive E° shows the reverse: \( \ce{Ag^+} \) is reduced more easily, so Ag is oxidised less easily than \( \ce{H2} \). This is why silver does not displace hydrogen from dilute acids.

• Option (4) \(Ag^+/Ag\) will act as anode when coupled with Standard Hydrogen electrode to form a cell — Incorrect. The electrode with the higher reduction potential is the cathode; here the SHE is the anode, as shown above.

NCERT (2026-27 edition):
Measurement of Electrode Potential – “… The standard electrode potentials are very important and we can extract a lot of useful information from them. The values of standard electrode potentials for some selected half-cell reduction reactions are given in Table 2.1. If the standard electrode potential of an electrode is greater than zero then its reduced form is more stable compared to hydrogen gas. Similarly, if the standard electrode potential is negative then hydrogen gas is more stable than the reduced form of the species.”

“A positive E^o means that the redox couple is a weaker reducing agent than the H⁺/H₂ couple.”