The equilibrium that exists in aqueous solution ,`CH_3COOHhArrCH_3COO^(-)+H^(+)` if dil HCI is added at constant temperature then

Options

A. The equilibrium constant will increase
B. The equilibrium constant will decrease
C. Concentration of `CH_3COO^(-)` will decrease
D. Concentration of `CH_3COO^(-)` will increase

Answer

To solve the problem regarding the equilibrium of acetic acid in the presence of dilute HCl, we will follow these steps:

### Step-by-Step Solution:

1. **Understand the Equilibrium Reaction**:
The equilibrium reaction given is:
\[ \text{CH}_3\text{COOH} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}^+ \]
This shows that acetic acid (a weak acid) dissociates into acetate ions and hydrogen ions.

2. **Effect of Adding HCl**:
When dilute HCl is added to the solution, it dissociates completely into H\(^+\) and Cl\(^-\) ions:
\[ \text{HCl} \rightarrow \text{H}^+ + \text{Cl}^- \]
Since HCl is a strong acid, it provides a significant amount of H\(^+\) ions to the solution.

3. **Le Chatelier's Principle**:
According to Le Chatelier's principle, if a system at equilibrium is disturbed, the system will shift in the direction that counteracts the disturbance. In this case, the increase in H\(^+\) concentration due to the addition of HCl will shift the equilibrium to the left (towards the reactants).

4. **Resulting Changes in Concentration**:
As the equilibrium shifts to the left, the concentration of acetic acid (CH\(_3\)COOH) will increase while the concentration of acetate ions (CH\(_3\)COO\(^-\)) will decrease. This is because the added H\(^+\) ions from HCl suppress the dissociation of acetic acid.

5. **Conclusion**:
Therefore, the concentration of acetate ions (CH\(_3\)COO\(^-\)) will decrease when dilute HCl is added to the solution.

### Final Answer:
The concentration of CH\(_3\)COO\(^-\) will **decrease** when dilute HCl is added.

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