Problem 44 - 43 Years JEE ADVANCED (1978-2020) + JEE MAIN Chapterwise & Topicwise Solved Papers Chemistry

Question

For the reduction of \(\mathrm{NO}_{3}^{-}\)ion in an aqueous solution, \(E^{\circ}\) is \(+0.96 \mathrm{~V}\). Values of \(E^{\circ}\) for some metal ions are given below \(\mathrm{V}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightarrow \mathrm{V}\) \(E^{\circ}=-1.19 \mathrm{~V}\) \(\mathrm{Fe}^{3+}(\mathrm{aq})+3 \mathrm{e}^{-} \rightarrow \mathrm{Fe} \quad E^{\circ}=-0.04 \mathrm{~V}\) \(\mathrm{Au}^{3+}(\mathrm{aq})+3 \mathrm{e}^{-} \rightarrow \mathrm{Au}\) \(E^{\circ}=+1.40 \mathrm{~V}\) \(\mathrm{Hg}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightarrow \mathrm{Hg} \quad E^{\circ}=+0.86 \mathrm{~V}\) The pair(s) of metals that is (are) oxidized by \(\mathrm{NQ}^{-}\)in aqueous solution is (are) [2009] (a) \(\mathrm{V}\) and \(\mathrm{Hg}\) (b) \(\mathrm{Hg}\) and \(\mathrm{Fe}\) (c) \(\mathrm{Fe}\) and \(\mathrm{Au}\) (d) Fe and \(\mathrm{V}\)

Step-by-Step Solution

Answer

The pair of metals oxidized by \(\mathrm{NO}_3^-\) are \(\mathrm{Fe}\) and \(\mathrm{V}\), option (d).

Step 1: Understanding Potentials

For a redox reaction, the substance with the higher standard reduction potential acts as an oxidizing agent, causing the reduction of the other substance with a lower potential, which acts as a reducing agent. The standard reduction potential of \(\mathrm{NO}_3^-\) (\(E^{\circ} = +0.96\, \mathrm{V}\)) indicates that anything with a potential less than \(+0.96\, \mathrm{V}\) can be oxidized by \(\mathrm{NO}_3^-\).

Step 2: Comparing Reduction Potentials

We compare the standard reduction potentials of the metals with that of \(\mathrm{NO}_3^-\). Those with potential less than \(+0.96\, \mathrm{V}\) will be oxidized by \(\mathrm{NO}_3^-\), as \(\mathrm{NO}_3^-\) will preferentially reduce. The reduction potentials of the metals are: \(\mathrm{V}: -1.19\, \mathrm{V}\), \(\mathrm{Fe}: -0.04\, \mathrm{V}\), \(\mathrm{Au}: +1.40\, \mathrm{V}\), \(\mathrm{Hg}: +0.86\, \mathrm{V}\).

Step 3: Identifying the Oxidized Metals

From the potentials, \(\mathrm{Hg}\) and \(\mathrm{Fe}\) have potentials of \(+0.86 \, \mathrm{V}\) and \(-0.04 \, \mathrm{V}\) which are both less than \(+0.96\, \mathrm{V}\). This means \(\mathrm{Hg}\) and \(\mathrm{Fe}\) can be oxidized by \(\mathrm{NO}_3^-\). The other metals, \(\mathrm{V} (-1.19 \, \mathrm{V})\) can also be oxidized, whereas \(\mathrm{Au}\, (+1.40\, \mathrm{V})\) cannot since its potential is higher.

Step 4: Conclusion

Based on the potentials, the metals that will be oxidized by \(\mathrm{NO}_3^-\) are \(\mathrm{Fe}\) and \(\mathrm{V}\). This corresponds to option (d).

Key Concepts

Standard Reduction Potential

Standard reduction potential is a crucial concept when discussing redox reactions. It helps us predict how a redox reaction will proceed by comparing the ease with which different substances gain electrons, or in other words, get reduced.
The standard reduction potential, represented by the symbol \(E^{\circ}\), is measured in volts and typically provided under standard conditions. These include a concentration of 1 M, a pressure of 1 atm, and a temperature of 25°C.
A positive \(E^{\circ}\) value indicates that the substance has a stronger tendency to gain electrons and acts as a good oxidizing agent. Conversely, a more negative \(E^{\circ}\) signifies a lesser tendency to gain electrons, implying the substance is more easily oxidized, often acting as a reducing agent.
By comparing \(E^{\circ}\) values, we can determine which species will be reduced and which will be oxidized in a given reaction. The substance with a higher reduction potential will act as the oxidizing agent and cause the substance with a lower potential to be oxidized.

Oxidizing Agent

An oxidizing agent is a substance that facilitates the oxidation of another substance while itself being reduced. It effectively "takes" electrons from the substance it oxidizes.
In a redox reaction, the oxidizing agent always pairs with a reducing agent, participating in the exchange of electrons.
When selecting an oxidizing agent in a chemical reaction, it's crucial to consider the standard reduction potentials (\(E^{\circ}\)) of the substances involved.

• The higher the \(E^{\circ}\) value, the better the substance acts as an oxidizing agent.
• The oxidizing agent will always have a more positive \(E^{\circ}\) compared to the substance it oxidizes.
• The oxidizing agent is reduced, meaning it gains electrons in the redox process.

For example, in the provided exercise, \(\mathrm{NO}_3^-\) with a \(E^{\circ} = +0.96\, \text{V}\), acts as an oxidizing agent, reducing other metals with lower \(E^{\circ}\) values.

Reducing Agent

The reducing agent is the opposite player in a redox reaction, donating electrons to another substance and causing it to be reduced. In doing so, the reducing agent itself gets oxidized.
Understanding the role of reducing agents is essential for predicting the direction and outcome of redox reactions, explained by comparing standard reduction potentials.
Here are some key points:

• The reducing agent has a lower \(E^{\circ}\) value compared to the oxidizing agent.
• It loses electrons during the reaction and is thus oxidized.
• It causes the oxidation of the other substance by donating electrons.

For instance, in the example from the exercise, \(\mathrm{V}\), \(\mathrm{Fe}\), and \(\mathrm{Hg}\), all have \(E^{\circ}\) values less than \(\mathrm{NO}_3^-\), making them candidates for reducing agents when reacting with \(\mathrm{NO}_3^-\).

Aqueous Solution Chemistry

In the world of chemistry, aqueous solutions are those where water is the solvent. These solutions are significant because many chemical reactions, including redox reactions, occur in aqueous mediums.
The behavior of solutes in water can greatly affect reaction dynamics and outcomes. When dealing with ions in aqueous solutions, like the metal ions mentioned in the exercise, several aspects come into play:

• Water facilitates the mobility of ions, which is crucial for electron transfer in redox reactions.
• It can also influence the standard reduction potential of the solute ions.
• The pH of the aqueous environment can alter reaction pathways and equilibrium.

Understanding how substances behave in aqueous solution chemistry is vital for accurately predicting and manipulating the behavior of ionic species during chemical reactions, such as the redox reactions highlighted in our example.