For an acidic indicator has Ka ​value of `10^(−5)`. The % basic form when pH equal to 5.3 is : [Given log2=0.3]

Options

A. 33.33
B. 66.66
C. 99
D. 11

Explanation

To solve the problem, we need to find the percentage of the basic form of an acidic indicator when the pH is 5.3, given that the \( K_a \) value is \( 10^{-5} \).

### Step-by-Step Solution:

1. **Calculate the concentration of \( H^+ \) ions:**
The pH is given as 5.3. We can use the formula:
\[ \text{pH} = -\log[H^+] \]
Rearranging this gives:
\[ [H^+] = 10^{-\text{pH}} = 10^{-5.3} \]
To compute \( 10^{-5.3} \):
\[ 10^{-5.3} = 10^{-5} \times 10^{-0.3} = 10^{-5} \times 0.5 \approx 5.011 \times 10^{-6} \, \text{M} \]

2. **Determine the degree of dissociation (\( \alpha \)):**
The degree of dissociation can be calculated using the formula:
\[ \alpha = \frac{K_a}{C} \]
where \( C \) is the concentration of the acid. Assuming the concentration of the acid is \( C = 5.011 \times 10^{-6} \, \text{M} \), we substitute the values:
\[ \alpha = \frac{10^{-5}}{5.011 \times 10^{-6}} \approx 1.995 \]

3. **Calculate the percentage of dissociation:**
The percentage of dissociation is given by:
\[ \text{Percentage of dissociation} = \frac{\alpha}{\alpha + 1} \times 100 \]
Substituting the value of \( \alpha \):
\[ \text{Percentage of dissociation} = \frac{1.995}{1.995 + 1} \times 100 \approx \frac{1.995}{2.995} \times 100 \approx 66.66\% \]

4. **Calculate the percentage of the basic form:**
Since the total percentage must equal 100%, the percentage of the basic form is:
\[ \text{Percentage of basic form} = 100\% - \text{Percentage of dissociation} \]
\[ \text{Percentage of basic form} = 100\% - 66.66\% \approx 33.34\% \]

### Final Answer:
The percentage of the basic form when pH equals 5.3 is approximately **33.34%**.