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Chapter 20: Problem 45

By using the data in Appendix E, determine whether each of the following substances is likely to serve as an oxidant or a reductant: (a) \(\mathrm{Cl}_{2}(g),(\mathbf{b}) \mathrm{MnO}_{4}^{-}(a q,\) acidic solution), (c) \(\mathrm{Ba}(s),(\mathbf{d}) \mathrm{Zn}(s) .\)

Short Answer

Expert verified
Based on the standard reduction potentials from Appendix E, the given substances are likely to act as follows: (a) \(\mathrm{Cl}_{2}(g)\) - Oxidant (b) \(\mathrm{MnO}_{4}^{-}(a q, \text {acidic solution})\) - Oxidant (c) \(\mathrm{Ba}(s)\) - Reductant (d) \(\mathrm{Zn}(s)\) - Reductant

Step by step solution

01

Find the standard reduction potentials for each substance

Consult Appendix E to find the standard reduction potential values (\(E^\circ\)) for each substance. Look for the half-reactions that directly involve the given substances. (a) For \(\mathrm{Cl}_{2}(g)\), find the half-reaction: \(\mathrm{Cl}_{2}(g)+2e^{-}\rightarrow 2 \mathrm{Cl}^{-}(a q)\). The \(E^\circ\) value for this half-reaction is 1.36 V. (b) For \(\mathrm{MnO}_{4}^{-}(a q, \text {acidic solution})\), find the half-reaction: \(\mathrm{MnO}_{4}^{-}+8 \mathrm{H}^{+}+5 e^{-}\rightarrow \mathrm{Mn}^{2+}+4 \mathrm{H}_{2} \mathrm{O}\). The \(E^\circ\) value for this half-reaction is 1.51 V. (c) For \(\mathrm{Ba}(s)\), find the half-reaction: \(\mathrm{Ba}^{2+}+2 e^{-}\rightarrow \mathrm{Ba}(s)\). The \(E^\circ\) value for this half-reaction is -2.92 V. (d) For \(\mathrm{Zn}(s)\), find the half-reaction: \(\mathrm{Zn}^{2+}+2 e^{-}\rightarrow \mathrm{Zn}(s)\). The \(E^\circ\) value for this half-reaction is -0.76 V.
02

Determine whether each substance acts as an oxidant or reductant

A substance will act as an oxidant if its reduction potential is high, indicating that it has a high tendency to gain electrons (undergo reduction). Conversely, a substance will act as a reductant if its reduction potential is low, indicating that it has a high tendency to lose electrons (undergo oxidation). (a) \(\mathrm{Cl}_{2}(g)\): Since the reduction potential is 1.36 V, a high positive value, \(\mathrm{Cl}_{2}(g)\) is likely to act as an oxidant. (b) \(\mathrm{MnO}_{4}^{-}(a q, \text {acidic solution})\): Since the reduction potential is 1.51 V, another high positive value, \(\mathrm{MnO}_{4}^{-}\) is also likely to act as an oxidant. (c) \(\mathrm{Ba}(s)\): With a reduction potential of -2.92 V, a low negative value, \(\mathrm{Ba}(s)\) is likely to act as a reductant. (d) \(\mathrm{Zn}(s)\): The reduction potential of -0.76 V, another low negative value, indicates that \(\mathrm{Zn}(s)\) is likely to act as a reductant.
03

Conclusion

Based on the data in Appendix E, the substances will likely act as oxidants or reductants as follows: (a) \(\mathrm{Cl}_{2}(g)\) - Oxidant (b) \(\mathrm{MnO}_{4}^{-}(a q, \text {acidic solution})\) - Oxidant (c) \(\mathrm{Ba}(s)\) - Reductant (d) \(\mathrm{Zn}(s)\) - Reductant

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Key Concepts

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

Standard Reduction Potentials
Standard reduction potential is a key factor in understanding redox reactions. It indicates the likelihood of a chemical species to be reduced, i.e., to gain electrons. Standard conditions, such as 25掳C, 1 atm pressure, and 1 M concentration, provide a reference point for these values. These potentials are measured in volts. When you look at standard reduction potentials:
  • Positive values mean a species readily gains electrons, thus acting as an oxidant.
  • Negative values imply the species is less willing to gain electrons, often acting as a reductant.
For instance, in our examples, - \(\mathrm{Cl}_{2}(g)\) and \(\mathrm{MnO}_{4}^{-}\) have positive potentials, meaning they are strong oxidants.- Conversely, \(\mathrm{Ba}(s)\) and \(\mathrm{Zn}(s)\) show negative potentials, classifying them as strong reductants. Understanding these potentials helps predict the direction of electron flow in chemical reactions.
Oxidants and Reductants
Oxidants and reductants are two primary components in redox reactions.The oxidant, or oxidizing agent, gains electrons. It facilitates oxidation by reducing itself in the process. Essentially, it 'pulls' electrons toward itself:
  • Oxidants have high standard reduction potentials.
  • They are often substances with nonmetallic characteristics like chlorine (\(\mathrm{Cl}_{2}\)).
Conversely, the reductant, or reducing agent, loses electrons. This process is accompanied by oxidation of the reductant:
  • Reductants show low standard reduction potentials.
  • They are usually metals such as barium (\(\mathrm{Ba}\)) and zinc (\(\mathrm{Zn}\)).
In a redox reaction, the oxidant and reductant work together to enable electron transfer, which is key to numerous chemical processes.
Electron Transfer
Electron transfer is a fundamental aspect of redox reactions. During these reactions, electrons are transferred from the reductant (which gets oxidized) to the oxidant (which gets reduced). This transfer is what drives the chemical reactions forward.
  • Electrons move from areas of lower potential (reductants) to higher potential (oxidants).
  • This movement is guided by the difference in standard reduction potentials.
The larger the difference in potentials, the greater the amount of energy that can be harnessed from the transfer. It's this energy from electron transfer that biological systems and batteries exploit. In our examples, when \(\mathrm{Zn}(s)\) gives up electrons, they are accepted by \(\mathrm{Cl}_{2}(g)\) or \(\mathrm{MnO}_{4}^{-}\) in suitable conditions, facilitating the whole redox process.

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Most popular questions from this chapter

A voltaic cell is constructed that is based on the following reaction: $$ \mathrm{Sn}^{2+}(a q)+\mathrm{Pb}(s) \longrightarrow \mathrm{Sn}(s)+\mathrm{Pb}^{2+}(a q) $$ (a) If the concentration of \(\mathrm{Sn}^{2+}\) in the cathode half-cell is 1.00\(M\) and the cell generates an emf of \(+0.22 \mathrm{V},\) what is the concentration of \(\mathrm{Pb}^{2+}\) in the anode half-cell? (b) If the anode half-cell contains \(\left[\mathrm{SO}_{4}^{2-}\right]=1.00 M\) in equilibrium with \(\mathrm{PbSO}_{4}(s),\) what is the \(K_{s p}\) of \(\mathrm{PbSO}_{4} ?\)

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?

In the Bronsted-Lowry concept of acids and bases, acid-base reactions are viewed as proton-transfer reactions. The stronger the acid, the weaker is its conjugate base. If we were to think of redox reactions in a similar way, what particle would be analogous to the proton? Would strong oxidizing agents be analogous to strong acids or strong bases? [Sections 20.1 and 20.2\(]\)

A voltaic cell similar to that shown in Figure 20.5 is constructed. One electrode half-cell consists of a silver strip placed in a solution of \(\mathrm{AgNO}_{3},\) and the other has an iron strip placed in a solution of \(\mathrm{FeCl}_{2}\) . The overall cell reaction is $$ \mathrm{Fe}(s)+2 \mathrm{Ag}^{+}(a q) \longrightarrow \mathrm{Fe}^{2+}(a q)+2 \mathrm{Ag}(s) $$ (a) What is being oxidized, and what is being reduced? (b) Write the half-reactions that occur in the two half-cells. (c) Which electrode is the anode, and which is the cathode? (d) Indicate the signs of the electrodes. (e) Do electrons flow from the silver electrode to the iron electrode or from the iron to the silver? (f) In which directions do the cations and anions migrate through the solution?

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. $$ \begin{array}{l}{\text { (a) } \mathrm{I}_{2} \mathrm{O}_{5}(s)+5 \mathrm{CO}(g) \longrightarrow \mathrm{I}_{2}(s)+5 \mathrm{CO}_{2}(g)} \\\ {\text { (b) } 2 \mathrm{Hg}^{2+}(a q)+\mathrm{N}_{2} \mathrm{H}_{4}(a q) \longrightarrow 2 \mathrm{Hg}(l)+\mathrm{N}_{2}(g)+4 \mathrm{H}^{+}(a q)} \\\ {\text { (c) } 3 \mathrm{H}_{2} \mathrm{S}(a q)+2 \mathrm{H}^{+}(a q)+2 \mathrm{NO}_{3}^{-}(a q) \longrightarrow 3 \mathrm{S}(s)+} \\\\{\quad\quad 2 \mathrm{NO}(g)+4 \mathrm{H}_{2} \mathrm{O}(l)}\end{array} $$
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