Assertion : The larger the value of equilibrium constant (K). Greater is the extent to which reactants are converted into the products.
Reason : A larger value of K facilitates the forward reactions to proceed to a greater extent before equilibrium is attained.

Options

A. If both Assertion and Reason are CORRECT and a Reason is the CORRECT explanations of the Assertion
B. If both Assertion and Reason are CORRECT but Reaon is not CORRECT explanation of the Assertion.
C. If Assertion is CORRECT but Reasson is INCORRECT.
D. If Assertion is INCORRECT but Reason is CORRECT.

Answer

To solve the question, we need to analyze both the assertion and the reason provided.

### Step-by-Step Solution:

1. **Understanding the Assertion**:
- The assertion states that "the larger the value of the equilibrium constant (K), the greater is the extent to which reactants are converted into products."
- This implies that a higher equilibrium constant indicates a greater tendency for the reaction to favor the formation of products.

2. **Understanding the Reason**:
- The reason states that "a larger value of K facilitates the forward reactions to proceed to a greater extent before equilibrium is attained."
- This means that when K is large, the forward reaction occurs more readily compared to the reverse reaction, allowing more products to form before reaching equilibrium.

3. **Relating K to Reaction Rates**:
- The equilibrium constant (K) is defined as the ratio of the concentrations of products to the concentrations of reactants at equilibrium.
- Mathematically, for a reaction \( aA + bB \rightleftharpoons cC + dD \):
\[ K = \frac{[C]^c [D]^d}{[A]^a [B]^b} \]
- If K is large, it indicates that the concentration of products ([C] and [D]) is much greater than that of reactants ([A] and [B]) at equilibrium.

4. **Forward and Backward Reaction Rates**:
- The rate of the forward reaction can be expressed as:
\[ \text{Rate}_{\text{forward}} = k_f [A]^a [B]^b \]
- The rate of the backward reaction can be expressed as:
\[ \text{Rate}_{\text{backward}} = k_b [C]^c [D]^d \]
- At equilibrium, these rates are equal:
\[ k_f [A]^a [B]^b = k_b [C]^c [D]^d \]

5. **Conclusion**:
- Since a larger K implies that \( k_f \) (the forward rate constant) is larger than \( k_b \) (the backward rate constant), it indicates that the forward reaction is favored, leading to more products being formed.
- Therefore, both the assertion and the reason are correct, and the reason correctly explains the assertion.

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

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