A $0.1\, M$ solution of $HF$ is $1\%$ ionized. What is the $K_a$
${10^{ - 5}}$
${10^{ - 4}}$
$3 \times {10^{ - 5}}$
$3 \times {10^{ - 4}}$
Given the two concentration of $HCN (K_a = 10^{-9})$ are $0.1\,M$ and $0.001\,M$ respectively. What will be the ratio of degree of dissociation ?
The first ionization constant of $H _{2} S$ is $9.1 \times 10^{-8}$. Calculate the concentration of $HS ^{-}$ ion in its $0.1 \,M$ solution. How will this concentration be affected if the solution is $0.1\, M$ in $HCl$ also? If the second dissociation constant of $H _{2} S$ is $1.2 \times 10^{-13}$, calculate the concentration of $S^{2-}$ under both conditions.
A weak base $MOH$ of $0.1\,N$ concentration shows a $pH$ value of $9$ . What is the percentage degree of ionization of the base ? .......$\%$
Ionic product of water at $310 \,K$ is $2.7 \times 10^{-14}$. What is the $\mathrm{pH}$ of neutral water at this temperature?
Which of the following will occur if a $0.1 \,M$ solution of a weak acid is diluted to $0.01\,M$ at constant temperature