Problem 67 - A Complete Resource Book in Chemistry for JEE Main

Question

AgCl precipitate dissolves in \(\mathrm{NH}_{4} \mathrm{OH}\) due to the formation of (a) \(\mathrm{Ag}\left(\mathrm{NH}_{3}\right)_{2} \mathrm{Cl}\) (b) \(\mathrm{Ag}\left(\mathrm{NH}_{3}\right)_{2} \mathrm{OH}\) (c) \(\mathrm{Ag}\left(\mathrm{NH}_{4}\right)_{2} \mathrm{Cl}\) (d) \(\mathrm{Ag}\left(\mathrm{NH}_{3}\right)_{2} \mathrm{Cl}_{2}\)

Step-by-Step Solution

Answer

The correct formation is (a) [Ag(NH_3)_2]Cl .

Step 1: Understand the Given Reaction

The given problem involves the dissolution of a silver chloride (AgCl) precipitate in an ammonium hydroxide ( NH_4OH ) solution. This process typically happens due to the formation of a soluble complex ion.

Step 2: Recognize Complex Ion Formation

When AgCl is dissolved in NH_4OH, it forms a complex ion. A common reaction for similar dissolutions involves silver ions (Ag^+) forming a complex with ammonia ( NH_3 ), such as the diammine silver complex ( [Ag(NH_3)_2]^+ ). In this case, silver ions will coordinate with ammonia molecules.

Step 3: Identify the Complex Formed

Given the options in the problem, only option (a) [Ag(NH_3)_2]Cl correctly suggests the formation of a complex ion with ammonia. This matches the typical dissolving pattern where AgCl in NH_4OH leads to Ag^+ ions coordinating with ammonia, forming [Ag(NH_3)_2]^+ ions which are soluble.

Step 4: Verify Through Reaction Equation

The reaction can be written as: AgCl(s) + 2NH_3(aq)  [Ag(NH_3)_2]^+ (aq) + Cl^- (aq) . This shows the formation of the diammine silver ion complex which makes AgCl dissolve due to its solubility in the presence of ammonia.

Key Concepts

Silver Chloride Dissolution

Silver chloride, AgCl, is known for its limited solubility in water. When mixed with water, only a tiny amount of it dissolves. However, its behavior changes significantly when introduced to certain chemicals, like ammonium hydroxide (NH_4OH).
When silver chloride comes into contact with ammonium hydroxide, it reacts to form new, soluble substances. The magic happens because of a key concept: complex ion formation. AgCl, which is typically insoluble, can dissolve in the presence of ammonia.
This occurs because the ammonia interacts with the silver ions, allowing AgCl to initially dissociate into silver ions (Ag^+) and chloride ions (Cl^-). Then, these silver ions team up with ammonia molecules, enabling AgCl to dissolve more than it would in plain water.

Diammine Silver Complex

The diammine silver complex, ":[Ag(NH_3)_2]^+", is the star player in the dissolution process of silver chloride in ammonia. In the reaction environment, Ag^+ ions are surrounded by and bond with ammonia molecules, forming this special complex.
This process is what allows AgCl to become soluble in the solution.
Here's a simple breakdown: the silver ion (Ag^+) doesn't stay alone. It finds ammonia molecules (NH_3) to associate with, creating a complex ion. This particular complex is stable and can remain dissolved in the solution.
Once formed, the diammine silver complex increases the overall solubility of silver salts in the mixture. This unique characteristic is widely utilized in chemistry to manipulate the solubility of silver compounds.

Ammonia Coordination Chemistry

Coordination chemistry revolves around how molecules like ammonia can bind to metals, such as silver. Ammonia is a simple molecule but can act as a ligand, meaning it can donate electrons to form bonds with metal ions.
In the case of silver chloride dissolution, ammonia plays a crucial role. When ammonia coordinates or sacrifices its electrons to interact with silver ions, it changes the dynamics of the solution, allowing typically insoluble compounds to dissolve.
This type of chemical interaction is described as a coordination bond, and in this case, it leads to the formation of the diammine silver complex, [Ag(NH_3)_2]^+. This bond makes the newly formed complexes stable and soluble.
Understanding coordination chemistry helps explain why some metals dissolve in certain solutions and not others. It also illustrates how specific ligand-metal interactions can dramatically alter the solubility and behavior of substances that would otherwise remain inert or insoluble.