From each of the following pairs of substances, use data in Appendix E to choose the one that is the stronger oxidizing agent: $$ \begin{array}{l}{\text { (a) } \mathrm{Cl}_{2}(g) \text { or } \mathrm{Br}_{2}(l)} \\ {\text { (b) } \mathrm{Zn}^{2+}(a q) \text { or } \mathrm{Cd}^{2+}(a q)} \\ {\text { (c) } \mathrm{Cl}^{-}(a q) \text { or } \mathrm{ClO}_{3}(a q)} \\ {\text { (d) } \mathrm{H}_{2} \mathrm{O}_{2}(a q) \text { or } \mathrm{O}_{3}(\mathrm{g})}\end{array} $$

To compare the oxidizing abilities of Cl₂(g) and Br₂(l), we will check their standard reduction potentials in Appendix E. The reaction for chlorine gas is: \[ \mathrm{Cl}_{2}(g) + 2 \mathrm{e}^{-} \rightarrow 2 \mathrm{Cl}^{-}(a q) \] The \( E° \) for this reaction is \( +1.36 \ \text{V} \). The reaction for bromine liquid is: \[ \mathrm{Br}_{2}(l) + 2 \mathrm{e}^{-} \rightarrow 2 \mathrm{Br}^{-}(a q) \] The \( E° \) for this reaction is \( +1.07 \ \text{V} \). Since Cl₂(g) has a higher standard reduction potential than Br₂(l), we can conclude that Cl₂(g) is the stronger oxidizing agent.

For this comparison, we will check the standard reduction potentials for Zn²⁺(aq) and Cd²⁺(aq). The reaction for zinc ions is: \[ \mathrm{Zn}^{2+}(a q) + 2 \mathrm{e}^{-} \rightarrow \mathrm{Zn}(s) \] The \( E° \) for this reaction is \( -0.76 \ \text{V} \). The reaction for cadmium ions is: \[ \mathrm{Cd}^{2+}(a q) + 2 \mathrm{e}^{-} \rightarrow \mathrm{Cd}(s) \] The \( E° \) for this reaction is \( -0.40 \ \text{V} \). Comparing the standard reduction potentials, Cd²⁺(aq) has a higher value than Zn²⁺(aq), which means Cd²⁺(aq) is the stronger oxidizing agent.

Let's compare the standard reduction potentials for Cl⁻(aq) and ClO₃⁻(aq). We already found the reaction for Cl⁻(aq) in part (a). However, oxidation is required in this case, so we need to reverse the reaction and change the sign of \( E° \): \[ \mathrm{2Cl}^{-}(a q) \rightarrow \mathrm{Cl}_{2}(g) + 2 \mathrm{e}^{-} \] with \( E° = -1.36 \ \text{V} \). The reaction for ClO₃⁻(aq) is: \[ \mathrm{2ClO}_{3}^{-}(a q) + 12 \mathrm{H}^{+}(a q) + 10 \mathrm{e}^{-} \rightarrow \mathrm{Cl}_{2}(g) + 6 \mathrm{H}_{2}\mathrm{O}(l) \] The \( E° \) for this reaction is \( +1.50 \ \text{V} \). Since ClO₃⁻(aq) has a higher standard reduction potential than Cl⁻(aq), we can conclude that ClO₃⁻(aq) is the stronger oxidizing agent.

Lastly, we will compare the standard reduction potentials for H₂O₂(aq) and O₃(g). The reaction for hydrogen peroxide is: \[ \mathrm{H}_{2} \mathrm{O}_{2}(a q) + 2 \mathrm{H}^{+}(a q) + 2 \mathrm{e}^{-} \rightarrow 2 \mathrm{H}_{2}\mathrm{O}(l) \] The \( E° \) for this reaction is \( +1.77 \ \text{V} \). The reaction for ozone is: \[ \mathrm{2O}_{3}(g) + 2 \mathrm{e}^{-} \rightarrow 3 \mathrm{O}_{2}(g) \] The \( E° \) for this reaction is \( +2.07 \ \text{V} \). Comparing the standard reduction potentials, O₃(g) has a higher value than H₂O₂(aq), so we can conclude that O₃(g) is the stronger oxidizing agent.