Question

Identify the substance oxidised reduced, oxidising agent and reducing agent for each of the following reactions: (a) \(2 \mathrm{AgBr}(\mathrm{s})+\mathrm{C}_{6} \mathrm{H}_{6} \mathrm{O}_{2}(\mathrm{aq}) \rightarrow 2 \mathrm{Ag}(\mathrm{s})+2 \mathrm{HBr}(\mathrm{aq})+\mathrm{C}_{6} \mathrm{H}_{4} \mathrm{O}_{2}(\mathrm{aq})\) (b) \(\begin{aligned} \mathrm{HCHO}(\mathrm{l})+2\left[\mathrm{Ag}\left(\mathrm{NH}_{3}\right)_{2}\right]^{+}(\mathrm{aq})+3 \mathrm{OH}^{-}(\mathrm{aq}) \rightarrow 2 \mathrm{Ag}(\mathrm{s})+\mathrm{HCOO}^{-}(\mathrm{aq})+4 \mathrm{NH}_{3} &(\mathrm{aq}) \\\\+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \end{aligned}\) (c) \(\mathrm{HCHO}(\mathrm{l})+2 \mathrm{Cu}^{2+}(\mathrm{aq})+5 \mathrm{OH}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cu}_{2} \mathrm{O}(\mathrm{s})+\mathrm{HCOO}^{-}(\mathrm{aq})+3 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})\) (d) \(\mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{l})+2 \mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{l}) \rightarrow \mathrm{N}_{2}(\mathrm{~g})+4 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})\) (e) \(\mathrm{Pb}(\mathrm{s})+\mathrm{PbO}_{2}(\mathrm{~s})+2 \mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}) \rightarrow 2 \mathrm{PbSO}_{4}(\mathrm{~s})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})\)

Short answer

(a) AgBr oxidizing agent, C6H6O2 reducing. (b) Silver complex oxidizing, HCHO reducing. (c) Cu2+ oxidizing, HCHO reducing. (d) H2O2 oxidizing, N2H4 reducing. (e) PbO2 oxidizing, Pb reducing.

Step-by-step solution

Analyzing Reaction (a)

In the reaction \(2 \mathrm{AgBr} + \mathrm{C}_{6} \mathrm{H}_{6} \mathrm{O}_{2} \rightarrow 2 \mathrm{Ag} + 2 \mathrm{HBr} + \mathrm{C}_{6} \mathrm{H}_{4} \mathrm{O}_{2}\), AgBr is reduced to Ag, and \(\mathrm{C}_{6} \mathrm{H}_{6} \mathrm{O}_{2}\) is oxidized to \(\mathrm{C}_{6} \mathrm{H}_{4} \mathrm{O}_{2}\). The silver ion (Ag+) in AgBr is the oxidizing agent, and the hydroxyl group of \(\mathrm{C}_{6} \mathrm{H}_{6} \mathrm{O}_{2}\) is the reducing agent.

Analyzing Reaction (b)

In the reaction \(\mathrm{HCHO} + 2[\mathrm{Ag}(\mathrm{NH}_{3})_{2}]^{+} + 3 \mathrm{OH}^{-} \rightarrow 2 \mathrm{Ag} + \mathrm{HCOO}^{-} + 4 \mathrm{NH}_{3} + 2 \mathrm{H}_{2} \mathrm{O}\), silver ions (from \(\left[\mathrm{Ag}(\mathrm{NH}_{3})_{2}\right]^{+})\) are reduced to silver metal (Ag), while the \(\mathrm{HCHO}\) is oxidized to \(\mathrm{HCOO}^{-}\). The silver complex is the oxidizing agent, and formaldehyde (\(\mathrm{HCHO}\)) is the reducing agent.

Analyzing Reaction (c)

For \(\mathrm{HCHO} + 2 \mathrm{Cu}^{2+} + 5 \mathrm{OH}^{-} \rightarrow \mathrm{Cu}_{2} \mathrm{O} + \mathrm{HCOO}^{-} + 3 \mathrm{H}_{2} \mathrm{O}\), Cu(II) ions are reduced to \(\mathrm{Cu}_{2} \mathrm{O}\), whereas \(\mathrm{HCHO}\) is oxidized to \(\mathrm{HCOO}^{-}\). Thus, \(\mathrm{Cu}^{2+}\) acts as the oxidizing agent while formaldehyde acts as the reducing agent.

Analyzing Reaction (d)

In the reaction \(\mathrm{N}_{2} \mathrm{H}_{4} + 2 \mathrm{H}_{2} \mathrm{O}_{2} \rightarrow \mathrm{N}_{2} + 4 \mathrm{H}_{2} \mathrm{O}\), \(\mathrm{N}_{2} \mathrm{H}_{4}\) is oxidized to \(\mathrm{N}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}_{2}\) is reduced to \(\mathrm{H}_{2} \mathrm{O}\). Therefore, hydrogen peroxide is the oxidizing agent, and hydrazine is the reducing agent.

Analyzing Reaction (e)

For \( \mathrm{Pb} + \mathrm{PbO}_{2} + 2 \mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow 2 \mathrm{PbSO}_{4} + 2 \mathrm{H}_{2} \mathrm{O}\), \(\mathrm{PbO}_{2}\) is reduced to \(\mathrm{PbSO}_{4}\) and \(\mathrm{Pb}\) is oxidized to \(\mathrm{PbSO}_{4}\). Thus, \(\mathrm{PbO}_{2}\) is the oxidizing agent and lead metal (\(\mathrm{Pb}\)) is the reducing agent.

Key concepts

Oxidation

Oxidation is a fundamental concept in redox reactions. When a substance undergoes oxidation, it loses electrons. This process results in an increase in the oxidation state of the element involved. For example:- In the reaction \[2 \text{AgBr} + \text{C}_6\text{H}_6\text{O}_2 \rightarrow 2 \text{Ag} + 2 \text{HBr} + \text{C}_6\text{H}_4\text{O}_2\]\(\text{C}_6\text{H}_6\text{O}_2\) is oxidized to \(\text{C}_6\text{H}_4\text{O}_2\).Here, the molecule \(\text{C}_6\text{H}_6\text{O}_2\) loses hydrogens (and therefore electrons) and gains an oxygen, portraying the classical characteristics of oxidation. This change indicates that the oxidation state of the hydroxyl group has increased.

Reduction

Reduction is the opposite of oxidation in redox reactions. It involves the gain of electrons by a substance. During reduction, the oxidation state of the element involved decreases. Consider this reaction:- In \[ \text{HCHO} + 2[\text{Ag}(\text{NH}_3)_2]^+ + 3 \text{OH}^- \rightarrow 2 \text{Ag} + \text{HCOO}^- + 4 \text{NH}_3 + 2 \text{H}_2\text{O}\] Silver ions (\([\text{Ag}(\text{NH}_3)_2]^+\)) are reduced to silver metal \((\text{Ag})\).The silver ions gain electrons during the reaction, resulting in the formation of solid silver. This electron gain embodies the reduction process, as the silver ions transition to a lower oxidation state by becoming neutral atoms.

Oxidizing Agent

An oxidizing agent, also known as an oxidant, is a substance that facilitates the oxidation of another substance. It does this by accepting electrons from the other substance and thereby gets reduced itself.Take for instance:- In the reaction\[ \text{HCHO} + 2 \text{Cu}^{2+} + 5 \text{OH}^- \rightarrow \text{Cu}_2\text{O} + \text{HCOO}^- + 3 \text{H}_2\text{O}\] The \(\text{Cu}^{2+}\) ions act as the oxidizing agent by oxidizing formaldehyde (\(\text{HCHO}\)) to formate ion (\(\text{HCOO}^-\)).Since the \(\text{Cu}^{2+}\) ions accept electrons from \(\text{HCHO}\), they effectively carry out oxidation on another substance by undergoing reduction themselves.

Reducing Agent

A reducing agent, or reductant, is a substance that causes the reduction of another substance. Reducing agents donate electrons to other substances and are thereby oxidized.Consider the reaction:- In\[ \text{N}_2\text{H}_4 + 2 \text{H}_2\text{O}_2 \rightarrow \text{N}_2 + 4 \text{H}_2\text{O}\] The \(\text{N}_2\text{H}_4\) acts as the reducing agent by donating electrons to \(\text{H}_2\text{O}_2\), which leads to its conversion to water.By donating electrons, \(\text{N}_2\text{H}_4\) facilitates the reduction of hydrogen peroxide (\(\text{H}_2\text{O}_2\)) to water (\(\text{H}_2\text{O}\)). Meanwhile, \(\text{N}_2\text{H}_4\) is oxidized to \(\text{N}_2\), demonstrating a classic redox process where one molecule serves as the reducing agent by providing electrons to another.