To address the question of what happens when a large quantity of NaCl is added to the aqueous solution of sodium stearate, let's apply Le Chatelier's principle, which helps us predict the direction in which an equilibrium will shift when a change occurs in the system.
### Step-by-Step Analysis:
1. Understand the Equilibrium:
[tex]\[
C_{17}H_{35}COONa (aq) \leftrightarrow C_{17}H_{35}COO^- (aq) + Na^+ (aq)
\][/tex]
2. Initial Condition:
- On the left side, we have sodium stearate (C[tex]\( _{17} \)[/tex]H[tex]\( _{35} \)[/tex]COONa).
- On the right side, we have stearate ions (C[tex]\( _{17} \)[/tex]H[tex]\( _{35} \)[/tex]COO[tex]\(^-\)[/tex]) and sodium ions (Na[tex]\(^+\)[/tex]).
3. Effect of Adding NaCl:
- When NaCl is added to the solution, it dissociates completely into Na[tex]\(^+\)[/tex] and Cl[tex]\(^-\)[/tex] ions in the solution.
- This increases the concentration of Na[tex]\(^+\)[/tex] ions in the solution.
4. Application of Le Chatelier's Principle:
- According to Le Chatelier's principle, if the concentration of one of the products in a reaction increases, the equilibrium will shift to counteract this change by reducing the concentration of that product.
- In this case, the increase in Na[tex]\(^+\)[/tex] ions will cause the equilibrium to shift to the left.
5. Shift to the Left:
- The equilibrium shifts to the left to reduce the concentration of Na[tex]\(^+\)[/tex] ions.
- This means more sodium stearate (soap) will form from the stearate ions and Na[tex]\(^+\)[/tex] ions.
6. Resulting Consequence:
- The shift to the left could lead to the precipitation of solid soap as the solution becomes supersaturated with respect to sodium stearate.
### Conclusion:
When a large quantity of NaCl is added to the aqueous solution, the equilibrium shifts to the left to such an extent that it leads to the precipitation of solid soap.
Thus, the correct answer is:
[tex]\[ \text{The equilibrium will shift to the left to such an extent that it leads to the precipitation of solid soap.} \][/tex]
