Problem 152 - A Complete Resource Book in Chemistry for JEE Main

Question

When \(\mathrm{NH}_{4} \mathrm{Cl}\) is added to an aqueous solution of \(\mathrm{NH}_{4} \mathrm{OH}\), (a) Conc. of \(\left[\mathrm{OH}^{-}\right]\)ions decreases. (b) Conc. of \(\left[\mathrm{OH}^{-}\right]\)ions increases. (c) Conc. of \(\left[\mathrm{NH}_{4}^{+}\right]\)ions as well as conc. \(\left[\mathrm{OH}^{-}\right]\)ions increase. (d) Conc. of \(\left[\mathrm{NH}_{4}^{+}\right]\)ions decreases.

Step-by-Step Solution

Answer

(a) The concentration of \([\mathrm{OH}^-]\) ions decreases.

Step 1: Understanding the Chemical Reaction

In an aqueous solution of \(\mathrm{NH}_4\mathrm{OH}\), \(\mathrm{NH}_4\mathrm{OH}\) partially dissociates into \(\mathrm{NH}_4^+\) and \(\mathrm{OH}^-\) ions due to its weak base nature. The equilibrium can be represented as:\[\mathrm{NH}_4\mathrm{OH} \rightleftharpoons \mathrm{NH}_4^+ + \mathrm{OH}^-\] Adding \(\mathrm{NH}_4\mathrm{Cl}\), which dissociates completely, will increase the concentration of \(\mathrm{NH}_4^+\) ions in the solution.

Step 2: Applying the Common Ion Effect

The increase in concentration of \(\mathrm{NH}_4^+\) ions due to the addition of \(\mathrm{NH}_4\mathrm{Cl}\) will shift the equilibrium of the dissociation of \(\mathrm{NH}_4\mathrm{OH}\) to the left according to Le Chatelier's principle. This means there will be a decrease in the concentration of \(\mathrm{OH}^-\) ions in the solution.

Step 3: Analyzing the Ion Concentrations

Since the equilibrium shifts to the left, the \(\mathrm{OH}^-\) ion concentration will decrease, while the \(\mathrm{NH}_4^+\) concentration is higher initially due to \(\mathrm{NH}_4\mathrm{Cl}\) addition. Therefore, the statement that matches this scenario is that the concentration of \([\mathrm{OH}^-]\) ions decreases.

Key Concepts

Le Chatelier's principle

Le Chatelier's principle is a fascinating concept in chemistry that helps us predict how changing conditions affect chemical equilibria. Imagine a group of acrobats perfectly balanced on a see-saw. If someone tips the see-saw to one side, the acrobats will shift to maintain their balance. Similarly, this principle states that if a stress is applied to a system at equilibrium, the system will shift in a direction that counteracts this stress, striving to restore equilibrium. When additional \(\mathrm{NH}_4^{+}\) ions from \(\mathrm{NH}_4 ext{Cl}\) are added to the solution, the reaction shifts to consume some of these excess ions, thereby decreasing the concentration of \([\mathrm{OH}^-\)] ions.

• This shift to the left (towards Reactants) occurs because the equilibrium tries to minimize the effect of the added common ion.
• In our exercise, it reduces the hydroxide ion concentration while effectively rebalancing the system.

Understanding this principle is crucial in predicting the impact of changing concentrations in reactions involving equilibria. It is like chemistry's way of keeping its balance, just like those acrobats.

Chemical equilibrium

Chemical equilibrium is a state within a chemical reaction where the concentrations of reactants and products cease to change. At equilibrium, the rate at which the reactants convert into products is equal to the rate at which products convert back into reactants. This dynamic balance is vital as it helps in understanding reaction properties and controlling industrial processes.
When \(\mathrm{NH}_4 ext{OH}\) dissociates, it establishes a chemical equilibrium with its ions:\[\mathrm{NH}_4 ext{OH} \rightleftharpoons \mathrm{NH}_4^+ + \mathrm{OH}^-\]

• The state of equilibrium signifies that, though reactions are still occurring, there is no net change in concentration.
• This balance is sensitive to outside influences, such as the addition of \(\mathrm{NH}_4 ext{Cl}\), which shifts the equilibrium, showing how chemicals can interact in a closed loop of reactions.

The understanding of equilibrium is paramount in chemistry, allowing researchers and students alike to predict how modifying reaction conditions can affect the outcome of a chemical system.

Ion concentration changes

The concentration of ions in a solution is a key factor in determining the behavior and properties of the solution. When another compound is added to a solution at equilibrium, concentrations of ions can change significantly due to the common ion effect.

• The common ion effect occurs when a salt containing an ion already present in the solution is added, affecting the solubility of compounds and the equilibrium state. For instance, adding \(\mathrm{NH}_4 ext{Cl}\) increases \(\mathrm{NH}_4^+\) ions.
• This addition can lead to a decrease in \([\mathrm{OH}^-\)] ion concentration due to a shift in equilibrium, based on Le Chatelier's principle.

Such changes in ion concentration have practical implications: they can alter pH levels, influence reaction rates, or affect the color of the solution. By understanding how ion concentrations fluctuate in response to common ions, chemists can manipulate reactions and solutions more effectively, which is pivotal in fields such as pharmaceuticals, environmental science, and industrial chemistry.