Question

From each of the following pairs of substances, use data in Appendix \(\mathrm{E}\) to choose the one that is the stronger reducing agent: (a) \(\mathrm{Al}(s)\) or \(\mathrm{Mg}(s)\) (b) \(\mathrm{Fe}(s)\) or \(\mathrm{Ni}(s)\) (c) \(\mathrm{H}_{2}(g\), acidic solution) or \(\operatorname{Sn}(s)\) (d) \(\mathrm{I}^{-}(a q)\) or \(\mathrm{Br}^{-}(a q)\)

Short answer

Based on the reduction potentials in Appendix E, the stronger reducing agents among the given pairs are: (a) Mg(s) with \(E^{\circ}=-2.37V\), (b) Fe(s) with \(E^{\circ}=-0.44V\), (c) Sn(s) with \(E^{\circ}=-0.14V\), and (d) I⁻(aq) with \(E^{\circ}=0.54V\).

Step-by-step solution

(a): Al(s) or Mg(s)

: 1. The reduction potentials are: Al³⁺ + 3e⁻ → Al(s): \(E^{\circ}=-1.66V\) Mg²⁺ + 2e⁻ → Mg(s): \(E^{\circ}=-2.37V\) 2. Compare the reduction potentials. Since -2.37V < -1.66V, the magnesium reduction potential is more negative. 3. The stronger reducing agent is Mg(s).

(b): Fe(s) or Ni(s)

: 1. The reduction potentials are: Fe²⁺ + 2e⁻ → Fe(s): \(E^{\circ}=-0.44V\) Ni²⁺ + 2e⁻ → Ni(s): \(E^{\circ}=-0.25V\) 2. Compare the reduction potentials. Since -0.44V < -0.25V, the iron reduction potential is more negative. 3. The stronger reducing agent is Fe(s).

(c): H2(g, acidic solution) or Sn(s)

: 1. The reduction potentials are: 2H⁺ + 2e⁻ → H₂: \(E^{\circ}=0.00V\) Sn²⁺ + 2e⁻ → Sn(s): \(E^{\circ}=-0.14V\) 2. Compare the reduction potentials. Since -0.14V < 0.00V, the tin reduction potential is more negative. 3. The stronger reducing agent is Sn(s).

(d): I⁻(aq) or Br⁻(aq)

: 1. The reduction potentials are: I₂ + 2e⁻ → 2I⁻: \(E^{\circ}=0.54V\) Br₂ + 2e⁻ → 2Br⁻: \(E^{\circ}=1.07V\) 2. Compare the reduction potentials. Since 0.54V < 1.07V, the iodine reduction potential is more negative. 3. The stronger reducing agent is I⁻(aq).

Key concepts

Reduction Potential

The reduction potential is a measure of the tendency of a chemical species to acquire electrons and thereby be reduced. The more negative or less positive the reduction potential, the greater the tendency of the species to donate electrons. This is crucial in determining the reactivity of different elements and compounds in redox reactions.

For instance, in the original exercise, the comparison between the reduction potentials of various substances, such as

• Aluminum: Al³⁺ + 3e⁻ → Al(s) with (-1.66V)
• Magnesium: Mg²⁺ + 2e⁻ → Mg(s) with (-2.37V)

shows that magnesium has a more negative reduction potential. This implies that magnesium has a greater tendency to lose electrons compared to aluminum. Hence, it serves as a stronger reducing agent. Calculating and comparing reduction potentials enables chemists to foresee the direction of electron flow in electrochemical cells and understand the spontaneity of reactions.

Reducing Agent

A reducing agent is a substance that donates electrons to another substance, causing itself to oxidize in the process. In simpler terms, a reducing agent loses electrons and facilitates the gain of electrons (or reduction) for another species. The effectiveness of a reducing agent is often determined by its reduction potential. When you pick

• Iron: Fe(s)

over

• Nickel: Ni(s)

as the stronger reducing agent, this decision is made by looking at their reduction potentials. Iron, with a potential of (-0.44V), donates electrons more readily than nickel (-0.25V). This concept is key in processes such as electrolysis, where specific reducing agents are chosen based on their ability to efficiently donate electrons. Strong reducing agents are essential for extracting metals from ores, synthesizing new compounds, and sustaining industrial processes.

Chemical Reactions

Chemical reactions involve the transformation of substances through breaking and forming of bonds. In the context of electrochemistry, reactions that involve the transfer of electrons are called redox reactions, which include oxidation and reduction processes. When evaluating redox reactions, knowing the roles of reactants, such as reducing agents, and their associated reduction potentials is important. Take for instance, when you compare the reaction of

• Hydrogen: 2H⁺ + 2e⁻ → H₂
• Tin: Sn²⁺ + 2e⁻ → Sn(s)

Reducing agents like tin are identified due to their more negative reduction potentials (-0.14V compared to 0.00V for hydrogen). This implies tin more readily gives up electrons to facilitate the electrochemical reactions.

Understanding how to balance redox reactions and predict their outcomes underpins a wide range of applications, from developing batteries and creating new materials to biological processes and environmental systems.