An indicator is a weak acid and $pH$ range is $4.0$ to $6.0$. If the indicator is $50\text{\%}$ ionized in a given solution then what is the ionization constant of the acid?

Q. The ionization constant of an acid-base indicator (a weak acid) is $1.0\times {10}^{-6}$. The ionized form of the indicator is red whereas the unionized form is blue. The $pH$ change required to alter the colour of the indicator from $80$% blue to $80$% red is:

Q. What is the $pH$ of a solution of $0.28M$ acid and $0.84M$ of its conjugate base if the ionization constant of acid is $4\times {10}^{-4}$?

Q. The ionization constant of a certain weak is ${10}^{-4}.$ What be the [salt] to [acid] ratio if we have to prepare a buffer with pH = 5 using the acid and one of the salts?

Q. In an acidic indicator HIn has ionization constant is ${10}^{-8}$. The acid form of indicator is yellow and alkaline form is red. Which is correct statement?

Given: $log\left(2\right)=0.3;log\left(3\right)=0.48$

Q. Calculate the pH range in which an acid indicator with $K_{acid}$(indicator) $=1.0\times {10}^{-5}$ changes colour when the concentration of the indicator is $1\times {10}^{-3}$ M.