Consider the following table of standard electrode potentials for a series of hypothetical reactions in aqueous solution: $$ \begin{array}{lr} \hline \text { Reduction Half-Reaction } & {E^{\circ}(\mathrm{V})} \\ \hline \mathrm{A}^{+}(a q)+\mathrm{e}^{-} \longrightarrow \mathrm{A}(s) & 1.33 \\\ \mathrm{~B}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{B}(s) & 0.87 \\\ \mathrm{C}^{3+}(a q)+\mathrm{e}^{-} \longrightarrow \mathrm{C}^{2+}(a q) & -0.12 \\ \mathrm{D}^{3+}(a q)+3 \mathrm{e}^{-} \longrightarrow \mathrm{D}(s) & -1.59 \\\ \hline \end{array} $$ (a) Which substance is the strongest oxidizing agent? Which is weakest? (b) Which substance is the strongest reducing agent? Which is weakest? (c) Which substance(s) can oxidize \(\mathrm{C}^{2+} ?\)

Step by step solution

01

a) Strongest and Weakest oxidizing agents

To determine the strongest and weakest oxidizing agents, we need to look at the electrode potentials in the given table since a more positive potential indicates a stronger oxidizing agent. The oxidation agents are the ones in the reduced form in the table. So, - Strongest oxidizing agent: A+, with an electrode potential of 1.33 V. - Weakest oxidizing agent: D3+, with an electrode potential of -1.59 V.

02

b) Strongest and Weakest reducing agents

To determine the strongest and weakest reducing agents, we need to look at the electrode potentials in the table. A more negative potential indicates a stronger reducing agent. The reducing agents are represented by the substances in their oxidized forms. To find the reducing agents, we need to change the half-reactions into their "opposite" form: 1. A(s) ⟶ A+(aq) + e⁻, E = -1.33 V 2. B(s) ⟶ B2+(aq) + 2e⁻, E = -0.87 V 3. C2+(aq) ⟶ C3+(aq) + e⁻, E = 0.12 V 4. D(s) ⟶ D3+(aq) + 3e⁻, E = 1.59 V So, - Strongest reducing agent: D(s), with reverted electrode potential: 1.59 V. - Weakest reducing agent: C2+ with reverted electrode potential: 0.12 V.

03

c) Substances that can oxidize C2+

Oxidizing C2+ means we need to find the substances that, when coupled with C2+, will have a positive overall cell potential (ΔE > 0). In other words, C2+ needs to be the strongest reducing agent among the species being compared. From the calculations in part (b), we already know that C2+ has a reverted electrode potential of 0.12 V. For any species S with E(S), we need to check if E(S) + E(C2+) > 0. Comparing with the other reverted electrode potentials: 1. A(s), E = -1.33 V: -1.33 + 0.12 = -1.21 V 2. B(s), E = -0.87 V: -0.87 + 0.12 = -0.75 V 3. D(s), E = 1.59 V: 1.59 + 0.12 = 1.71 V The overall cell potential is positive only for D(s). So, only D(s) can oxidize C2+.

Unlock Step-by-Step Solutions & Ace Your Exams!

• Full Textbook Solutions

  Get detailed explanations and key concepts

• Unlimited Al creation

  Al flashcards, explanations, exams and more...

• Ads-free access

  To over 500 millions flashcards

• Money-back guarantee

  We refund you if you fail your exam.

Start your free trial

Over 30 million students worldwide already upgrade their learning with Vaia!