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Chapter 20: Problem 6

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 \(\mathrm{C}^{2+} ?\)

Short Answer

Expert verified
(a) Strongest oxidizing agent: A+, with an electrode potential of 1.33 V. Weakest oxidizing agent: D3+, with an electrode potential of -1.59 V. (b) Strongest reducing agent: D(s), with reverted electrode potential: 1.59 V. Weakest reducing agent: C2+ with reverted electrode potential: 0.12 V. (c) Only D(s) can oxidize C2+.

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) 鉄� A+(aq) + e鈦�, E = -1.33 V 2. B(s) 鉄� B2+(aq) + 2e鈦�, E = -0.87 V 3. C2+(aq) 鉄� C3+(aq) + e鈦�, E = 0.12 V 4. D(s) 鉄� D3+(aq) + 3e鈦�, 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 (螖E > 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+.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Understanding Oxidizing Agents
Oxidizing agents are substances that gain electrons during a chemical reaction. This action of gaining electrons causes them to oxidize another substance, hence the name "oxidizing agents." They are crucial in redox reactions, where one substance is oxidized, and another is reduced. The potential of an oxidizing agent to obtain electrons is indicated by its electrode potential value 鈥� the more positive this value, the stronger the oxidizing power. In our exercise, the strongest oxidizing agent is \( A^+ \), with a potential of 1.33 V, because it can readily accept electrons. On the flip side, the weakest oxidizing agent is \( D^{3+} \), at -1.59 V, indicating reluctance to gain electrons.
Examining Reducing Agents
Reducing agents are chemicals that donate electrons to another molecule in a redox reaction. When they donate electrons, they become oxidized themselves, but in this process, they reduce the other substance, earning their title "reducing agents." To identify the strength of a reducing agent, we must consider the reversed electrode potential from the standard table. Substances with more negative potentials make stronger reducing agents, showing they have a greater tendency to lose electrons. From the given table after reversing the reactions, \( D(s) \) is determined to be the strongest reducing agent with a potential of -1.59 V, and \( C^{2+} \) is the weakest with 0.12 V.
Decoding Redox Reactions
Redox reactions involve the transfer of electrons between two substances. This amalgamation of oxidation and reduction processes happens simultaneously, as one molecule's loss of electrons (oxidation) directly corresponds to another's electron gain (reduction). Redox reactions are pivotal in both natural and technological processes, from cellular respiration in our bodies to energy generation in batteries. In our exercise, to determine which substances can oxidize \( C^{2+} \), we need to compare the electrode potentials of other candidates in reversed reactions with \( C^{2+} \)'s potential. Here, \( D(s) \) with a potential of -1.59 V in its reversed form is the only substance capable of oxidizing \( C^{2+} \), as it meets the criterion of a positive overall cell potential.

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Most popular questions from this chapter

For a spontaneous reaction \(\mathrm{A}(a q)+\mathrm{B}(a q) \longrightarrow \mathrm{A}^{-}(a q)+\) \(\mathrm{B}^{+}(a q),\) answer the following questions: (a) If you made a voltaic cell out of this reaction, what halfreaction would be occurring at the cathode, and what half reaction would be occurring at the anode? (b) Which half-reaction from (a) is higher in potential energy? (c) What is the sign of \(E_{\text {cell }}^{\circ}\) ?

(a) What is electrolysis? (b) Are electrolysis reactions thermodynamically spontaneous? (c) What process occurs at the anode in the electrolysis of molten \(\mathrm{NaCl}\) ? (d) Why is sodium metal not obtained when an aqueous solution of \(\mathrm{NaCl}\) undergoes electrolysis?

For each of the following balanced oxidation-reduction reactions, (i) identify the oxidation numbers for all the elements in the reactants and products and (ii) state the total number of electrons transferred in each reaction. (a) \(\mathrm{H}_{2}(g)+\mathrm{F}_{2}(g) \longrightarrow 2 \mathrm{HF}(g)\) (b) \(2 \mathrm{Fe}^{2+}(a q)+\mathrm{H}_{2} \mathrm{O}_{2}(a q)+2 \mathrm{H}^{+}(a q) \longrightarrow 2 \mathrm{Fe}^{3+}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\) (c) \(\mathrm{CH}_{4}(g)+2 \mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l)\)

Indicate whether each statement is true or false: (a) The anode is the electrode at which oxidation takes place. (b) A voltaic cell always has a positive emf. (c) A salt bridge or permeable barrier is necessary to allow a voltaic cell to operate.

Mercuric oxide dry-cell batteries are often used where a flat discharge voltage and long life are required, such as in watches and cameras. The two half-cell reactions that occur in the battery are $$ \begin{array}{l} \mathrm{HgO}(s)+\mathrm{H}_{2} \mathrm{O}(l)+2 \mathrm{e}^{-} \longrightarrow \mathrm{Hg}(l)+2 \mathrm{OH}^{-}(a q) \\ \mathrm{Zn}(s)+2 \mathrm{OH}^{-}(a q) \longrightarrow \mathrm{ZnO}(s)+\mathrm{H}_{2} \mathrm{O}(l)+2 \mathrm{e}^{-} \end{array} $$ (a) Write the overall cell reaction. (b) The value of \(E_{\text {red }}^{\circ}\) for the cathode reaction is \(+0.098 \mathrm{~V}\). The overall cell potential is \(+1.35 \mathrm{~V}\). Assuming that both half-cells operate under standard conditions, what is the standard reduction potential for the anode reaction? (c) Why is the potential of the anode reaction different than would be expected if the reaction occurred in an acidic medium?
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