Equilibrium practice · Mixed difficulty

Ionisation of weak acid can be calculated by the formula

Consider the following gaseous equilibrium with equilibrium constants ${K_1}$ and ${K_2}$ respectively
$\begin{array}{l} S{O_2}(g) + \frac{1}{2}{O_2}(g)\underset{{}}{\overset{{}}{\leftrightharpoons}}S{O_3}(g)\\ 2S{O_3}(g)\underset{{}}{\overset{{}}{\leftrightharpoons}}2S{O_2}(g) + {O_2}(g) \end{array}$
The equilibrium constants are related as

In the hydrolysis of a salt of weak acid and weak base, the hydrolysis constant ${K_h}$ is equal to

Partial pressure of $A,B,C$ and $D$ on the basis of gaseous system, $A + 2B\;\leftrightharpoons\;C + 3D,$ are $A = 0.20,\;B = 0.10,\;C = 0.30$ and $D = 0.50\;\mathrm{atm}$. The numerical value of equilibrium constant is

For a reaction in equilibrium :

$0.15$ mole of pyridinium chloride has been added into $500 \mathrm{~mL}$ of $0.2 \mathrm{~M}$ pyridine solution. Find the $\mathrm{pH}$ and $\left[\mathrm{OH}^{-}\right]$ ion concentration in the resulting solution. $\mathrm{K}_{\mathrm{b}}=1.5 \times 10^{-9} \mathrm{M}$ (assuming no change in volume).

For the reaction, ${\mathrm{H}_2}\left( \mathrm{g} \right) + {\mathrm{I}_2}\left( \mathrm{g} \right) \to 2\mathrm{HI}\left( \mathrm{g} \right)$, the equilibrium constant ${K_p}$ changes with

${N}_2 {O}_4(\mathrm{g}) \rightleftharpoons 2 {NO}_2(\mathrm{g})$

For the reaction, if the percentage dissociation of ${N}_2 {O}_4$ are $25 \%, 50 \%, 75 \%$ and $100 \%$, then the sequence of observed vapour densities will be $\left({d}_1, {~d}_2, {~d}_3\right.$ and ${d}_4$ are vapour density at $25 \%, 50 \%$, $75 \%$ and $100 \%$ dissociation respectively)

Which of the following has lowest ${P^H}$ value

The $pH$ of tears coming out of a person’s eye is:

The stronger Bronsted base is:

Which of the following is least likely to behave as Lewis acid?

For the gaseous reaction, ${\mathrm{C}_2}{\mathrm{H}_4} + {\mathrm{H}_2}\leftrightharpoons{\mathrm{C}_2}{\mathrm{H}_6},\;\Delta H = - 130\;\mathrm{kJ}\;\mathrm{ mo}{\mathrm{l}^{ - 1}}$ carried in a closed vessel, the equilibrium concentration of the ${\mathrm{C}_2}{\mathrm{H}_6}$ can definitely be increased by

Given, $p K_b$ values
$\begin{array}{cccccccccccccccccccc} {{\text{Species}}}&{p{K_b}} \\ {N{H_3}}&{4.74} \\ {{A^ - }}&{4.00} \\ {{B^ - }}&{5.00} \\ {{C^ - }}&{6.00} \end{array}$
queous solutions of $0.01 \mathrm{M}$ salts of $\mathrm{HX}$

$(X=A, B, C)$ and $\mathrm{NH}_4 \mathrm{OH}$ have $\mathrm{pH}$ in the order

A buffer solution is $0.25 \mathrm{~M}-\mathrm{CH}_3 \mathrm{COOH}+0.15 \mathrm{~M}-\mathrm{CH}_3 \mathrm{COONa},$ saturated in $\mathrm{H}_2 \mathrm{S}(0.1 \mathrm{~M})$ and has $\left[\mathrm{Mn}^{2+}\right]=0.04 \mathrm{~M}, K_{\mathrm{a}}\left(\mathrm{CH}_3 \mathrm{COOH}\right)=2.0 \times 10^{-5},$ $K_{\mathrm{a}}\left(\mathrm{H}_2 \mathrm{S}\right)=1.0 \times 10^{-21}$ and $K_{\mathrm{sp}}(\mathrm{MnS})=2.5 \times 10^{-13}.$
Which buffer component should be increased in concentration and to which minimum value to just start precipitation of $\mathrm{MnS}?$

For the reaction equilibrium,

$2\mathrm{NOBr}\left( \mathrm{g} \right)\;\leftrightharpoons\;2\mathrm{NO}\left( \mathrm{g} \right) + \mathrm{B}{\mathrm{r}_2}\left( \mathrm{g} \right),$ if ${P_{\mathrm{B}{\mathrm{r}_2}}} = \frac{P}{9}$ at equilibrium and $P$ is total pressure. The ratio ${K_p}/P$ is equal to:

Which of the following salts is the most basic in aqueous solution?

If the equilibrium constant for ${\mathrm{N}_\mathrm{2}}\left( \mathrm{g} \right)\,\mathrm{ + }\,{\mathrm{O}_\mathrm{2}}\left( \mathrm{g} \right)\,\,\,\leftrightharpoons\,\,\,\mathrm{2NO}\left( \mathrm{g} \right)\,$ is k, the equilibrium constant for $\frac{\mathrm{1}}{\mathrm{2}}{\mathrm{N}_\mathrm{2}}\left( \mathrm{g} \right)\,\mathrm{ + }\,\frac{\mathrm{1}}{\mathrm{2}}\,{\mathrm{O}_\mathrm{2}}\left( \mathrm{g} \right)\,\,\leftrightharpoons\,\,\mathrm{NO}\left( \mathrm{g} \right)$ will be:

Calculate the partial pressure of carbon monoxide from the following datas

$\mathrm{CaC}{\mathrm{O}_3}\mathrm{C}\xrightarrow{\Delta }\mathrm{CaO}\left( \mathrm{g} \right) + \mathrm{C}{\mathrm{O}_2}\; \uparrow \;{K_p} = 8 \times {10^{ - 2}}$

${\mathrm{O}_2}\left( \mathrm{g} \right) + \mathrm{C}\left( s \right) \to 2\mathrm{CO}\left( \mathrm{g} \right),\;\;\;{K_p} = 2$

Statement I :$\mathrm{SO}_{2(g)}+\frac{1}{2} \mathrm{O}_{2(g)} \rightleftharpoons S \mathrm{O}_{2(g)}$+heat. Forward rection is favoured at high temperature and low pressure.

Statement II : Reaction is endothermic.

A solution of $\mathrm{FeC}{\mathrm{l}_3}$ in water acts as acidic due to:

The first and second dissociation constants of an acid
${\mathrm{H}_2}A\,\mathrm{are}\,\,1.0 \times {10^{ - 5}}\,\mathrm{and}\,\,5.0 \times {10^{ - 10}}$ respectively. The overall dissociation constant of the acid will be

If $K_c$ is in the range of _____ appreciable concentrations of both reactants and products are present.

Which of the following will have maximum pH?

$\mathrm{20}\;\mathrm{mL}$ of a $0.1\;N\;\mathrm{HCl}$ is mixed with $\mathrm{20}\;\mathrm{mL}$ of a $0.1\;N\;\mathrm{KOH}$ solution. The $\mathrm{pH}$ of the solution would be:

For the reaction $C(s) + C{O_2}(g)2\,\,CO(g)$ the partial pressure of $C{O_2}$ and are 2.0 and 4.0 atm resp, at equilibrium. The ${K_p}$ for the reaction is

In which of the following reactions, increases in the volume at constant temperature do not affect the number of moles at equilibrium?

$pH$ of $0.05\;\;M\;\mathrm{Mg}{\left( {\mathrm{OH}} \right)_2}$ is:

Which of the following equilibria will shift to right side on increasing the temperature

At $600\,,\,{K_p}$ for the following reaction is 1 atm.

$X\left( \mathrm{g} \right)\,\,\,\leftrightharpoons\,\,\,Y\left( \mathrm{g} \right) + Z\left( \mathrm{g} \right)$

At equilibrium, $50\,\%$ of $X\left( \mathrm{g} \right)$ is dissociated. The total pressure of the equilibrium system is $p$ atm. What is the partial pressure (in atm) of $X\left( \mathrm{g} \right)$ at equilibrium?