Question

In the following reactions, label the oxidizing agent and the reducing agent. (a) $$ 2 \mathrm{Al}(s)+3 \mathrm{~F}_{2}(g) \longrightarrow 2 \mathrm{AlF}_{3}(s) $$ (b) $$ \begin{aligned} \mathrm{Hg}^{2+}(a q)+\mathrm{NO}_{2}^{-}(a q)+& \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \\ & \mathrm{Hg}(s)+2 \mathrm{H}^{+}(a q)+\mathrm{NO}_{3}^{-}(a q) \end{aligned} $$

Short answer

(a) Oxidizing agent: \(\mathrm{F}_{2}\), Reducing agent: \(\mathrm{Al}\); (b) Oxidizing agent: \(\mathrm{Hg}^{2+}\), Reducing agent: \(\mathrm{NO}_{2}^{-}\).

Step-by-step solution

Understand Redox Reactions

In redox (reduction-oxidation) reactions, oxidation is the process by which an atom, ion, or molecule loses electrons, while reduction is the process by which an atom, ion, or molecule gains electrons. The oxidizing agent is the substance that gets reduced by gaining electrons, and the reducing agent is the substance that gets oxidized by losing electrons.

Identify Reactants and Products (a)

In reaction (a): \[2 \mathrm{Al}(s) + 3 \mathrm{F}_{2}(g) \rightarrow 2 \mathrm{AlF}_{3}(s)\]Aluminum (Al) loses electrons to form \(\mathrm{Al^{3+}}\) ions in aluminum fluoride (\(\mathrm{AlF}_{3}\)), and fluorine (\(\mathrm{F}_{2}\)) gains electrons to form \(\mathrm{F^{-}}\) ions.

Determine Oxidation and Reduction Processes (a)

Aluminum (Al) goes from a 0 oxidation state to a +3 state, so it is oxidized and is the reducing agent. Fluorine \((\mathrm{F}_{2})\) goes from a 0 oxidation state to a -1 state, so it is reduced and is the oxidizing agent.

Identify Reactants and Products (b)

In reaction (b): \[\mathrm{Hg^{2+}}(aq) + \mathrm{NO_{2}^{-}}(aq) + \mathrm{H_{2}O}(l) \rightarrow \mathrm{Hg}(s) + 2 \mathrm{H^{+}}(aq) + \mathrm{NO_{3}^{-}}(aq)\]Mercury (Hg) ions gain electrons to form solid mercury, and nitrite ions (\(\mathrm{NO_{2}^{-}}\)) are oxidized to form nitrate ions (\(\mathrm{NO_{3}^{-}}\)).

Determine Oxidation and Reduction Processes (b)

Mercury ions (\(\mathrm{Hg^{2+}}\)) go from a +2 to a 0 oxidation state, so they are reduced and the oxidizing agent. The nitrite ion (\(\mathrm{NO_{2}^{-}}\)) goes from a +3 to a +5 oxidation state, so it is oxidized and is the reducing agent.

Key concepts

Oxidizing Agent

In chemistry, an oxidizing agent plays a crucial role in oxidation-reduction reactions, often abbreviated as redox reactions. It is the substance that gets reduced by accepting electrons from another species. This results in the oxidizing agent being reduced while the other substance, often called the reducing agent, loses electrons.

Key points about oxidizing agents:

• They are typically rich in oxygen or are nonmetals capable of accepting electrons.
• When they gain electrons, their oxidation state decreases.

Let's consider reaction (a) from the original exercise. Here, fluorine ( F_{2}) is the oxidizing agent. It gains electrons from aluminum (Al), going from a 0 oxidation state to a -1 oxidation state.

In reaction (b), mercury ions ( Hg^{2+}) act as the oxidizing agent. These ions gain electrons and are reduced to neutral mercury (Hg) in the process. By accepting electrons, oxidizing agents drive the oxidation of other species, enabling the reduction of their own state.

Reducing Agent

The reducing agent is crucial in redox reactions, often providing the electrons needed for reduction. This agent is oxidized itself, meaning it loses electrons during the chemical process. The nature of the reducing agent is to donate electrons, thereby increasing its own oxidation state while decreasing that of the oxidizing agent.

Key points about reducing agents:

• They are typically metals or rich in hydrogen, capable of donating electrons.
• Their oxidation state increases as they lose electrons.

In exercise reaction (a), aluminum (Al) serves as the reducing agent. Aluminum loses three electrons, transitioning from an oxidation state of 0 to +3.

For reaction (b), the nitrite ion ( NO_{2}^{-}) functions as the reducing agent. It loses electrons and is oxidized to nitrate ( NO_{3}^{-}), thus facilitating the reduction of another species.

Electron Transfer

At the heart of any redox reaction is the transfer of electrons, a fundamental process that leads to changes in the oxidation states of the elements involved. This electron transfer is significant because it underpins the entire redox process, allowing the roles of oxidizing and reducing agents to be fulfilled.

Important aspects of electron transfer in redox reactions:

• The substance that loses electrons becomes oxidized.
• The substance that gains these electrons is reduced.
• Electrons are often transferred in pairs, especially in reactions involving covalent compounds.

In the context of exercise reaction (a), electrons are transferred from aluminum to fluorine. This transfer allows aluminum to oxidize and fluorine to reduce.

Similarly, in reaction (b), electrons are transferred from the nitrite ion ( NO_{2}^{-}) to mercury ions ( Hg^{2+}). Each time electrons are handed over, a change in the chemical nature and oxidation states of the substances is observed. Understanding this flow of electrons is critical for mastering redox reactions.