Problem 121 - A Complete Resource Book in Chemistry for JEE Main

Question

Zn gives \(\mathrm{H}_{2}\) gas with \(\mathrm{H}_{2} \mathrm{SO}_{4}\) and \(\mathrm{HCl}\), but not wit \(\mathrm{HNO}_{3}\) because (a) in electrochemical series, \(\mathrm{Zn}\) is above hydrogen (b) \(\mathrm{NO}_{3}^{-}\)ion is reduced in preference to hydroniur ion (c) \(\mathrm{Zn}\) acts as an oxidizing agent when reacts wit \(\mathrm{HNO}_{3}\) (d) \(\mathrm{HNO}_{3}\) is weaker acid than \(\mathrm{H}_{2} \mathrm{SO}_{4}\) and \(\mathrm{HCl}\)

Step-by-Step Solution

Answer

The reason is that in the presence of HNO_3, NO_3^{-} ions are reduced in preference to hydronium ions (option b).

Step 1: Understanding the Role of Electrochemical Series

In the electrochemical series, elements such as Zn are above hydrogen. This means Zn can displace hydrogen ions from acids like HCl and H_2SO_4 to produce hydrogen gas. However, the activity series also determines the substances that will react with each other, impacting reactivity with certain acids.

Step 2: Analyzing Reaction with HNO_3

When Zn is placed in HNO_3, the typical displacement reaction that produces H_2 gas does not occur. Instead, a redox reaction takes place where the NO_3^{-} ion is preferentially reduced instead of the hydronium ion. This means the nitrate ion takes electrons rather than facilitating the release of hydrogen gas.

Step 3: Identifying Oxidizing and Reducing Agents

In the context of Zn's reaction with HNO_3, Zn loses electrons to the nitrates rather than the protons in the solution. This causes the metal to act as a reducing agent for the nitrate ion, implicating a redox process. It leads to the formation of nitrogen oxides rather than hydrogen gas.

Step 4: Evaluating Acid Strength

The suggestion that HNO_3 is weaker than H_2SO_4 and HCl is incorrect; HNO_3 is actually a strong acid. The reason hydrogen gas is not liberated is due to the selective reduction of nitrate ions, not because of acid strength.

Key Concepts

Redox Reactions

Redox reactions are fascinating chemical processes involving the transfer of electrons between two substances. The term "redox" is a shorthand for reduction-oxidation, encapsulating two complementary actions: reduction is the gain of electrons, while oxidation is the loss.
A redox reaction can be thought of as a balancing act of electrons where the substance gaining electrons is reduced and the substance losing them is oxidized. For example, during the reaction of zinc (Zn) with nitric acid ( HNO_3}), zinc does not displace hydrogen ions to form hydrogen gas as it does with other acids like hydrochloric acid ( HCl}) or sulfuric acid ( H_2SO_4}). Instead, Zn undergoes oxidation, losing electrons to nitrate ions ( NO_3^-}). This action simultaneously reduces the nitrate ions, which is a hallmark of a redox reaction.
Thus, the usual gas-forming displacement doesn’t occur because both Zn and the nitrate ions engage in a redox process, intricately entwining their roles.

Displacement Reactions

Displacement reactions are intriguing types of reactions where one element is replaced by another in a compound. In the context of metals and acids, a more reactive metal can displace a less reactive one, or hydrogen in the case of acids.
This is governed by the electrochemical series—a ranking that tells us about the reactivity of different chemical elements. In the scenario where Zn reacts with HCl} or H_2SO_4}, Zn is able to displace hydrogen ions, releasing hydrogen gas as a result. This happens because Zn is more reactive than hydrogen, as indicated by its position above hydrogen in the electrochemical series.
However, when HNO_3} is introduced, displacement does not occur. Instead, the nitrate ions prefer to participate in a redox reaction, capturing electrons from Zn instead of permitting hydrogen liberation. This makes it clear that the nature of the reacting acid profoundly influences the outcome, going beyond a simple displacement consideration.

Oxidizing and Reducing Agents

In any redox reaction, there are two key players involved: the oxidizing agent and the reducing agent. Understanding their roles is essential to grasp the full picture of such reactions.
The oxidizing agent is the substance that gains electrons, thus being reduced itself. Conversely, the reducing agent is the one that donates electrons, consequently being oxidized. Considering the Zn and HNO_3} reaction, Zn serves as the reducing agent. It donates electrons to the nitrate ions, undergoing oxidation in the process. This electron transfer reduces the nitrate ions, which are the oxidizing agents in this case.
This distinction between the roles of Zn and nitrate underscores how their interaction deviates from simpler substitution reactions. Instead of displacing hydrogen, they engage in a redox reaction resulting in the formation of new chemical species such as nitrogen oxides, rather than hydrogen gas. Such reactions exemplify the intricacies of chemical interactions driven by electron transfer.