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 $\mathrm{pH}=-\log \left[\mathrm{H}_{3} \mathrm{O}^{+}\right]$.
Explain ionization and ionization constant in di and polyprotic acid.
The $pH$ of the solution obtained on neutralisation of $40\, mL\, 0.1\, M\, NaOH$ with $40\, mL\, 0.1\, M\, CH_3COOH$ is
Accumulation of lactic acid $(HC_3H_5O_3),$ a monobasic acid in tissues leads to pain and a feeling of fatigue. In a $0.10\, M$ aqueous solution, lactic acid is $3.7\%$ dissociates. The value of dissociation constant, $K_a,$ for this acid will be
The ionization constant of $0.1$ $M$ weak acid is $1.74 \times {10^{ - 5}}$ at $298$ $K$ temperature. Calculate $pH$ of its $0.1$ $M$ solution.
Assuming that the degree of hydrolysis is small, the $pH$ of $0.1\, M$ solution of sodium acetate $(K_a\, = 1.0\times10^{- 5})$ will be