Assertion : Solubility of `AgCl` in `NH_(3)(aq)` is greater than in pure water.
Reason : When `AgCl` dissolve in `NH_(3)(aq)`, complex ion [`Ag(NH_(3))_(2)^(+)`] formation takes place and solubility equilibrium of `AgCl_(3)` shifted in forward direction.

Options

A. If both the statements are TRUE and STATEMENT-2 is the correct explation of STATEMENT-1
B. If both the statements are TRUE AND STATEMENT-2 is NOT the correct explanation of STATEMENT-1
C. If STATEMENT-1 is TRUE and STATEMETN-2 is FLASE
D. If STATEMENT-1 is FLASE and STATEMENT-2 is TRUE

Answer

To solve the question regarding the solubility of AgCl in NH₃(aq) compared to pure water, we will analyze the assertion and reason step by step.

### Step-by-Step Solution:

1. **Understanding the Assertion**:
- The assertion states that the solubility of AgCl in NH₃(aq) is greater than in pure water.
- We need to consider the solubility product (Ksp) of AgCl and how it is affected by the presence of NH₃.

**Hint**: Think about how solubility is defined and what factors can influence it.

2. **Dissolution of AgCl in Water**:
- In pure water, AgCl dissociates according to the following equilibrium:
\[ \text{AgCl (s)} \rightleftharpoons \text{Ag}^+ (aq) + \text{Cl}^- (aq) \]
- The solubility product (Ksp) expression for AgCl can be written as:
\[ K_{sp} = [\text{Ag}^+][\text{Cl}^-] \]

**Hint**: Consider what happens to the concentration of ions when AgCl dissolves in water.

3. **Dissolution of AgCl in NH₃(aq)**:
- When AgCl is added to NH₃(aq), it reacts to form a complex ion:
\[ \text{Ag}^+ (aq) + 2 \text{NH}_3 (aq) \rightleftharpoons \text{Ag(NH}_3)_2^+ (aq) \]
- This reaction indicates that Ag⁺ ions are being stabilized by forming a complex with NH₃, which shifts the equilibrium to favor the dissolution of more AgCl.

**Hint**: Think about how the formation of a complex ion affects the equilibrium position.

4. **Effect of Complex Formation**:
- The formation of the complex ion \(\text{Ag(NH}_3)_2^+\) reduces the concentration of free Ag⁺ ions in solution.
- According to Le Chatelier's principle, if the concentration of one of the products (Ag⁺) decreases, the equilibrium will shift to the right to produce more Ag⁺ ions, thus increasing the solubility of AgCl.

**Hint**: Recall Le Chatelier's principle and how it applies to changes in concentration.

5. **Conclusion**:
- Since the formation of the complex ion increases the solubility of AgCl in NH₃(aq), the assertion is true.
- The reason provided correctly explains why the solubility increases due to complex ion formation, thus the reason is also true.

### Final Answer:
Both the assertion and reason are correct, and the reason correctly explains the assertion.