Examkin

Question

Difficulty: Very hardElectrochemical Series and Reaction Spontaneity
The standard reduction potentials for three half-cells are given below:
Cr3+(aq)+3e−→Cr(s)E∘=−0.74 V\text{Cr}^{3+}(aq) + 3e^- \rightarrow \text{Cr}(s) \quad E^\circ = -0.74\text{ V}
Fe2+(aq)+2e−→Fe(s)E∘=−0.44 V\text{Fe}^{2+}(aq) + 2e^- \rightarrow \text{Fe}(s) \quad E^\circ = -0.44\text{ V}
Ag+(aq)+e−→Ag(s)E∘=+0.80 V\text{Ag}^+(aq) + e^- \rightarrow \text{Ag}(s) \quad E^\circ = +0.80\text{ V}

A chemistry student intends to store an aqueous solution of iron(II) tetraoxonitrate(V), Fe(NO3)2\text{Fe(NO}_3)_2, in metal containers. Based on standard electrode potentials, which container choice is suitable for safely storing the solution without undergoing a spontaneous redox reaction?

  1. Only the Silver containerAnswer
  2. B
    Only the Chromium container
  3. C
    Both the Chromium and Silver containers
  4. D
    Neither container

Answer

Only the Silver container can safely store the iron(II) tetraoxonitrate(V) solution.
A redox reaction is spontaneous if the standard cell potential is positive (E°cell > 0). When storing aqueous iron(II) ions in a Silver container, the potential reaction involves oxidation of Silver metal and reduction of iron(II) ions: E°cell = E°(reduction) - E°(oxidation) = -0.44 V - (+0.80 V) = -1.24 V. Because E°cell is negative, the reaction cannot occur spontaneously, making the Silver container suitable. Conversely, for Chromium, E°cell = -0.44 V - (-0.74 V) = +0.30 V, which is positive and causes a spontaneous reaction that destroys the Chromium container.

Step-by-Step Solution

1
Determine the cell potential E°cell for storing iron(II) solution in a Chromium container.
E°cell = E°cathode - E°anode = E°(Fe²⁺/Fe) - E°(Cr³⁺/Cr) = -0.44 V - (-0.74 V) = +0.30 V.
Chromium acts as the anode (oxidation) and iron(II) as the cathode (reduction).
2
Evaluate spontaneity for the Chromium container reaction.
Since E°cell = +0.30 V > 0, the reaction is spontaneous.
A positive standard cell potential indicates a feasible redox reaction, meaning the Chromium container will dissolve and react with the solution.
3
Determine the cell potential E°cell for storing iron(II) solution in a Silver container.
E°cell = E°(Fe²⁺/Fe) - E°(Ag⁺/Ag) = -0.44 V - (+0.80 V) = -1.24 V.
Silver acts as the anode (oxidation) and iron(II) as the cathode (reduction).
4
Evaluate spontaneity for the Silver container reaction and conclude safe storage.
Since E°cell = -1.24 V < 0, the reaction is non-spontaneous, so the Silver container safely holds the solution.
A negative cell potential means no spontaneous reaction occurs between Silver metal and aqueous iron(II) ions.

Key Concept

Spontaneity of Redox Reactions and Standard Cell Potential
Rate this question

Topics

Electrochemical Series and Reaction Spontaneity

More questions from this topic