A weak acid-type indicator was found to be 60% dissociated at pH = 9.18. What will be the percentage — Ionic Equilibrium Chemistry Question

💡 Solution & Explanation

Using the indicator equation: $\text{pH} = pK_a + \log(\frac{\alpha}{1-\alpha})$. At 60% dissociation: $9.18 = pK_a + \log(\frac{0.6}{0.4}) = pK_a + \log(1.5)$. Since $\log(1.5) = \log 3 - \log 2 = 0.48 - 0.30 = 0.18$, we get $pK_a = 9.18 - 0.18 = 9.00$. When the pH is exactly 9.0 (which equals the $pK_a$), the logarithmic term must be 0, meaning the ratio $\frac{\alpha}{1-\alpha} = 1$. Thus, the indicator is 50% dissociated. Formatted to four digits, this is 0050.