Question

In which direction will the following reaction go if the standard reduction potentials are 0.80 V for Ag/Ag and –0.44 V for Fe/Fe2?

Option 1: Ag Fe → Ag Fe2
Option 2: Ag Fe → Fe Ag2
Option 3: Fe Ag → Ag Fe2
Option 4: Fe Ag → Fe Ag2

Answer 1

The redox reaction will proceed in the direction where iron is oxidized and silver is reduced, which, according to the standard reduction potentials provided, is Option 2: Ag Fe → Fe Ag2 (corrected to 2Ag+ + Fe(s) → 2Ag(s) + Fe2+(aq)).

Step-by-step explanation:

The question deals with the direction in which a redox reaction will proceed based on the given standard reduction potentials. We have been given two half-reactions with their respective potentials: the reduction of silver, Ag/Ag+, which has a more positive standard reduction potential of 0.80 V, and the oxidation of iron, Fe/Fe2+, which has a negative standard reduction potential of -0.44 V.

To find the direction in which the reaction will proceed, we can use the formula:

Since the more positive Ecathode corresponds to a reduction reaction and the more negative Eanode corresponds to an oxidation reaction, the silver reaction will act as the cathode (reduction), and the iron reaction will act as the anode (oxidation). The cell potential will be positive, indicating that the reaction is spontaneous. Hence, the reaction will proceed in the direction:

Fe(s) + 2Ag+(aq) → Fe2+(aq) + 2Ag(s)

The correct choice based on the options given is Option 2: Ag Fe → Fe Ag2, which should be corrected to 2Ag+ + Fe(s) → 2Ag(s) + Fe2+(aq).

Answer 2

The reaction where Ag+ is reduced to Ag and Fe is oxidized to Fe2+ (Option 2: Ag + Fe → Fe2+ + Ag) is the correct direction, as Ag/Ag+ has a higher standard reduction potential than Fe/Fe2+.

Step-by-step explanation:

The question asks in which direction a reaction involving silver and iron will proceed based on their standard reduction potentials. Silver (Ag/Ag+) has a standard reduction potential of 0.80 V and iron (Fe/Fe2+) has a standard reduction potential of –0.44 V. To determine the direction of the reaction, we compare these values. The species with the higher (more positive) reduction potential will undergo reduction, while the species with the lower reduction potential will be oxidized. Thus, Ag+ will be reduced to Ag, and Fe will be oxidized to Fe2+.

The correct option is 'Option 2: Ag + Fe → Fe2+ + Ag', which corresponds to the overall balanced reaction 2Ag+ (aq) + Fe(s) ⇒ 2Ag(s) + Fe2+ (aq). The silver ion, being more positive, tends to gain electrons and reduce to solid silver, while iron tends to lose electrons and oxidize to Fe2+.