Explain a general step-wise approach to evaluate the $pH$ of the weak electrolyte.
Step-$1$ : The species present before dissociation are identified as Bronsted-Lowry acid/base.
Step-$2$: Balanced equations for all possible reaction. i.e. with a species acting both as acid as well as base are written.
Step-$3$: The reaction with the higher $\mathrm{K}_{a}$ is identified as the primary reaction whilst the other is a subsidiary reaction.
Step-$4$: Enlist in a tabular form the following values for each of the species in the primary reaction. $(i)$ Initial concentration $c$, $(ii)$ Change in concentration on proceeding to equilibrium in term of $(\alpha)$, degree of ionization. $(iii)$ Equilibrium concentration.
Step-$5$ : Substitute equilibrium concentrations into equilibrium constant equation for principal reaction and solve for $\alpha$.
Step-$6$: Calculate the concentration of species in principal reaction.
Step-$7$: Calculate $\mathrm{pH}=-\log \left[\mathrm{H}_{3} \mathrm{O}^{+}\right]$.
Derive the equation of relation between weak base ionization constant ${K_b}$ and its conjugate acid ionization constant ${K_a}$
$K_b$ for $NH_4OH$ is $1.8\times 10^{-5}.$ The $[\mathop O\limits^\Theta H]$ of $0.1\,M\,NH_4OH$ is
The degree of dissociation of $0.1\,M\,HCN$ solution is $0.01\%$ . Its ionisation constant would be
A compound whose aqueous solution will have the highest $pH$
The first and second dissociation constants of an acid $H_2A$ are $1.0 \times 10^{-5}$ and $5.0 \times 10^{-10}$ respectively. The overall dissociation constant of the acid will be