Assertion : In a titration of weak monoprotic acid with strong base, the `pH` at the half equivalence point is `pK_(a)`.
Reason : At half equivalence point, it will form acidic buffer at its maximum capacity where [acid]`=`[salt].

Options

A. If both the statements are TRUE and STATEMENT-2 is the correct explanation 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 STATEMENT-2 is FALSE
D. If STATEMENT-1 is FALSE and STATEMENT-2 is TRUE

Answer

To solve the question, we need to analyze both the assertion and the reason provided, and then prove their validity step by step.

### Step-by-Step Solution:

1. **Understanding the Assertion**:
- The assertion states that in a titration of a weak monoprotic acid with a strong base, the pH at the half equivalence point is equal to the pKa of the acid.
- A weak monoprotic acid can be represented as HA, where it can donate one proton (H⁺).

**Hint**: Remember that the half equivalence point is where half of the acid has been neutralized by the base.

2. **Titration Reaction**:
- When a weak acid (HA) is titrated with a strong base (e.g., NaOH), the reaction can be written as:
\[ HA + OH^- \rightarrow A^- + H_2O \]
- Here, A⁻ is the conjugate base formed from the weak acid.

**Hint**: Identify the species involved in the reaction and their roles.

3. **Defining the Half Equivalence Point**:
- At the half equivalence point, the amount of weak acid (HA) is equal to the amount of its conjugate base (A⁻). Thus, we have:
\[ [HA] = [A^-] \]

**Hint**: The half equivalence point is crucial for understanding the buffer system.

4. **Buffer Solution Formation**:
- At this point, we have a buffer solution consisting of equal concentrations of the weak acid and its conjugate base. The pH of a buffer solution can be described by the Henderson-Hasselbalch equation:
\[ pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right) \]
- Since \([A^-] = [HA]\), substituting this into the equation gives:
\[ pH = pK_a + \log(1) = pK_a + 0 \]
- Therefore, we conclude that:
\[ pH = pK_a \]

**Hint**: Use the Henderson-Hasselbalch equation to relate pH and pKa in buffer solutions.

5. **Validating the Reason**:
- The reason states that at the half equivalence point, an acidic buffer is formed at its maximum capacity where \([acid] = [salt]\). This is indeed correct, as we have established that the concentrations of the weak acid and its conjugate base are equal.

**Hint**: Recognize that the formation of a buffer is key to understanding the pH relationship.

6. **Conclusion**:
- Both the assertion and the reason are correct. The reason provides a valid explanation for the assertion. Therefore, we conclude that:
- Both statements are true, and the reason is the correct explanation for the assertion.

**Final Answer**: Both the assertion and reason are true, and the reason is the correct explanation for the assertion.