Calculate the ${H^ + }$ ion concentration in a $1.00\,(M)$ $\,HCN\,$ litre solution $({K_a} = 4 \times {10^{ - 10}})$
$4 \times {10^{ - 14}}\,mole/litre$
$2 \times {10^{ - 5}}\,mole/litre$
$2.5 \times {10^{ - 5}}\,mole/litre$
None of these
What is the percent ionization $(\alpha)$ of a $0.01\, M\, HA$ solution ? .......$\%$ $(K_a = 10^{-6})$
Degree of dissociation of $0.1\,N\,\,C{H_3}COOH$ is (Dissociation constant $ = 1 \times {10^{ - 5}}$)
The $pH$ of $0.1$ $M$ solution of cyanic acid $(HCNO)$ is $2.34$. Calculate the ionization constant of the acid and its degree of ionization in the solution.
Concentration $C{N^ - }$ in $0.1\,M\,HCN$ is $[{K_a} = 4 \times {10^{ - 10}}]$
If $pK_a =\, -\,log K_a=4$ for a weak acid $HX$ and $K_a= C\alpha ^2$ then Van't Haff factor when $C = 0.01\,M$ is