$5.0$ $pH$ containing solution is dilute $100$ times. Calculate $pH$ of dilute solution.
$5.0$ $pH$ containing solution is dilute $100$ times. Calculate $pH$ of dilute solution.
$7.0$
Similar Questions
$pH$ of $0.1\,\, M$ $N{H_3}$ aqueous solution is $({K_b} = 1.8 \times {10^{ - 5}})$
$pH$ of $0.1\,\, M$ $N{H_3}$ aqueous solution is $({K_b} = 1.8 \times {10^{ - 5}})$
The solubility of a salt of weak acid $( A B )$ at $pH 3$ is $Y \times 10^{-3} mol L ^{-1}$. The value of $Y$ is
. . . . . (Given that the value of solubility product of $A B \left( K _{ sp }\right)=2 \times 10^{-10}$ and the value of ionization constant of $H B \left( K _{ a }\right)=1 \times 10^{-8}$ )
The solubility of a salt of weak acid $( A B )$ at $pH 3$ is $Y \times 10^{-3} mol L ^{-1}$. The value of $Y$ is
. . . . . (Given that the value of solubility product of $A B \left( K _{ sp }\right)=2 \times 10^{-10}$ and the value of ionization constant of $H B \left( K _{ a }\right)=1 \times 10^{-8}$ )
- [IIT 2018]
Derive ${K_w} = {K_a} \times {K_b}$ and ${K_w} = p{K_a} \times p{K_b}$ for weak base $B$ and its conjugate acid ${B{H^ + }}$.
Derive ${K_w} = {K_a} \times {K_b}$ and ${K_w} = p{K_a} \times p{K_b}$ for weak base $B$ and its conjugate acid ${B{H^ + }}$.
When $CO_2$ dissolves in water, the following equilibrium is established
$C{O_2} + 2{H_2}O\, \rightleftharpoons {H_3}{O^ + } + HCO_3^ - $
for which the equilibrium constant is $3.8 \times 10^{-7}$ and $pH = 6.0$. The ratio of $[HCO_3^- ]$ to $[CO_2]$ would be :-
When $CO_2$ dissolves in water, the following equilibrium is established
$C{O_2} + 2{H_2}O\, \rightleftharpoons {H_3}{O^ + } + HCO_3^ - $
for which the equilibrium constant is $3.8 \times 10^{-7}$ and $pH = 6.0$. The ratio of $[HCO_3^- ]$ to $[CO_2]$ would be :-
A weak monoprotic acid of $0.1\, M,$ ionizes to $1\% $ in solution. What will be the $pH $ of solution
A weak monoprotic acid of $0.1\, M,$ ionizes to $1\% $ in solution. What will be the $pH $ of solution