Explain a general step-wise approach to evaluate the $pH$ of a weak electrolyte.

(N/A) Step-$1$: Identify the species present before dissociation as Bronsted-Lowry acids or bases.
Step-$2$: Write balanced equations for all possible reactions,including species acting as both acids and bases.
Step-$3$: Identify the reaction with the higher $K_{a}$ (or $K_{b}$) as the primary reaction,while others are considered subsidiary reactions.
Step-$4$: Create an $ICE$ (Initial,Change,Equilibrium) table for the primary reaction: $(i)$ Initial concentration $c$,$(ii)$ Change in concentration at equilibrium in terms of $\alpha$ (degree of ionization),$(iii)$ Equilibrium concentration.
Step-$5$: Substitute equilibrium concentrations into the equilibrium constant expression for the principal reaction and solve for $\alpha$.
Step-$6$: Calculate the concentration of the relevant species ($[H_{3}O^{+}]$ or $[OH^{-}]$) from the principal reaction.
Step-$7$: Calculate $pH$ using the formula $pH = -\log [H_{3}O^{+}]$.