Application of Equilibrium Constant

When two reactants, A and B are mixed to give products C and D, the reaction quotient $\mathrm{Q}$, at the initial stage of the reaction :–

Consider three vessels containing aqueous solutions of a weak acid $\mathrm{HA}$ (initial concentration ${\mathrm{C}}_1$, degree of dissociation ${\mathrm{\alpha }}_1$ ), another weak acid $\mathrm{HB}$ (initial concentration ${\mathrm{C}}_2$, degree of dissociation ${\mathrm{\beta }}_2$ ), and their mixture (initial concentration of $\mathrm{HA}={\mathrm{C}}_1$, degree of dissociation ${\mathrm{\alpha }}_3$ and initial concentration of $\mathrm{HB}={\mathrm{C}}_2$, degree of dissociation ${\mathrm{\beta }}_3$ ), respectively. When ${\mathrm{C}}_1$ and ${\mathrm{C}}_2$ are varied, the correct plot (assuming $\mathrm{\alpha }$ and $\mathrm{\beta }$ to be small) is

${\mathrm{N}}_2{\mathrm{O}}_4\left(\mathrm{g}\right)\rightleftharpoons 2{\mathrm{NO}}_2\left(\mathrm{g}\right), {\mathrm{K}}_{\mathrm{c}}=4.$   

The above reaction is studied graphically as shown in the figure.

$$ Which of the following statements are correct?

${\mathrm{COCl}}_2$ gas dissociates according to the equation, $\mathrm{COC}{\mathrm{l}}_2\underset{}{\rightleftharpoons }\mathrm{CO}\left(\mathrm{g}\right)+{\mathrm{Cl}}_2\left(\mathrm{g}\right)$. When heated to $700 \mathrm{K}$ the density of the gas mixture at $1.16 \mathrm{atm}$ and at equilibrium is $1.16 \mathrm{g}/\mathrm{litre}$. The degree of dissociation of ${\mathrm{CO}}_2$ at $700 \mathrm{K}$ is

When the reaction $\mathrm{A}+2\mathrm{B}\rightleftharpoons 2\mathrm{C}+\mathrm{D}$ was studied, it was observed that the initial concentration of $\mathrm{B}$ was $1.5$ times that of $\mathrm{A},$ and the equilibrium concentrations of $\mathrm{A}$ and $\mathrm{C}$ were equal. Then ${\mathrm{K}}_{\mathrm{c}}$ for the given equilibrium equals

Identify the correct expression for the equilibrium constant of the following reaction:
$2{\mathrm{X}}_{\left(\mathrm{g}\right)}+{\mathrm{Y}}_{\left(\mathrm{g}\right)}\rightleftharpoons 3{\mathrm{Z}}_{\left(\mathrm{g}\right)}$

${\mathrm{N}}_2{\mathrm{O}}_4\left(\mathrm{g}\right)\rightleftharpoons 2{\mathrm{NO}}_2\left(\mathrm{g}\right); \mathrm{Kc}=5.7\times {10}^{-9}$ at $298 \mathrm{K}$. At equilibrium

If $\mathrm{Kc}$ is very large 

$0.6$ mole of ${\mathrm{P}\mathrm{C}\mathrm{l}}_5,\mathrm{ }0.3$ mole of ${\mathrm{P}\mathrm{C}\mathrm{l}}_3$ and $0.5$ mole of ${\mathrm{C}\mathrm{l}}_2$ are taken in a $1\mathrm{L}$ flask to obtain the following equilibrium:
$\mathrm{P}\mathrm{C}{\mathrm{l}}_{5\left(\mathrm{g}\right)}\leftrightharpoons \mathrm{P}\mathrm{C}{\mathrm{l}}_{3\left(\mathrm{g}\right)}+\mathrm{C}{\mathrm{l}}_{2\left(\mathrm{g}\right)}$
If the equilibrium constant ${\mathrm{K}}_{\mathrm{c}}$ for the reaction is $0.2$. Predict the direction of the reaction.

$2\mathrm{A}+\mathrm{B}\rightleftharpoons 3\mathrm{C}, {\mathrm{K}}_{\mathrm{C}}=49$

For the above reaction, $3$ $\mathrm{L}$ vessel contains $2, 1$ and $3$ moles of $\mathrm{A}, \mathrm{B}$ and $\mathrm{C}$ respectively. Then, what will happen to the reaction?

The reaction quotient $\left(\mathrm{Q}\right)$ predicts:

In an aqueous solution of volume $500 \mathrm{mL},$ when the reaction of $2{\mathrm{Ag}}^{+}+{\mathrm{Cu}}_{\left(\mathrm{s}\right)}\rightleftharpoons {\mathrm{Cu}}^{2+}+2{\mathrm{Ag}}_{\left(\mathrm{s}\right)}$ reached equilibrium the $\left[{\mathrm{Cu}}^{2+}\right]$ was $\mathrm{x}\hspace{0.17em}\mathrm{M}$. When $500 \mathrm{mL}$ of water is further added, at the new equilibrium $\left[{\mathrm{Cu}}^{+2}\right]$ will be :-

The equilibrium constant for the reaction between ${\mathrm{CH}}_4 \left(\mathrm{g}\right)$ and ${\mathrm{H}}_2\mathrm{S} \left(\mathrm{g}\right)$ to form ${\mathrm{CS}}_2 \left(\mathrm{g}\right)$ and ${\mathrm{H}}_2 \left(\mathrm{g}\right)$, at $1173 \mathrm{K}$ is $3.6$. For the following composition of the reaction mixture, decide which of the following option is correct?
$$\begin{gathered} \left[{\mathrm{CH}}_4\right]=1.07 \mathrm{M}, \left[{\mathrm{H}}_2\mathrm{S}\right]=1.20 \mathrm{M}, \\ \left[{\mathrm{CS}}_2\right]=0.90 \mathrm{M} , \left[{\mathrm{H}}_2\right]=1.78 \mathrm{M} \end{gathered}$$

A mixture of ${\mathrm{NO}}_2$ and ${\mathrm{N}}_2{\mathrm{O}}_4$ has a vapour density of $38.3$ at $300 \mathrm{K}$. What is the number of moles of ${\mathrm{NO}}_2$ in $100 \mathrm{g}$ of the mixture?

The equilibrium constant for the reaction between ${\mathrm{CH}}_4 \left(\mathrm{g}\right)$ and ${\mathrm{H}}_2\mathrm{S} \left(\mathrm{g}\right)$ to form ${\mathrm{CS}}_2 \left(\mathrm{g}\right)$ and ${\mathrm{H}}_2 \left(\mathrm{g}\right)$, at $1173 \mathrm{K}$ is $3.6$. For the following composition of the reaction mixture, decide which of the following option is correct?
$$\begin{gathered} \left[{\mathrm{CH}}_4\right]=1.07 \mathrm{M}, \left[{\mathrm{H}}_2\mathrm{S}\right]=1.20 \mathrm{M}, \\ \left[{\mathrm{CS}}_2\right]=0.90 \mathrm{M} , \left[{\mathrm{H}}_2\right]=1.78 \mathrm{M} \end{gathered}$$

The equilibrium constant for the reaction, $\mathrm{A}+2\mathrm{B}\rightleftharpoons \mathrm{C}+\mathrm{D},$ is $1.0\times {10}^8.$

Calculate the equilibrium concentration of $\mathrm{A}$, if $1.0$ mole of $\mathrm{A}$ and $3.0$ mole of $\mathrm{B}$ are placed in $1\mathrm{L}$ flask and allowed to attain the equilibrium.

The equilibrium constant(K) for the reaction Cyclohexane $\rightleftharpoons$ Hex-1-ene is 10732. At the given condition in a closed one litre vessel 2 moles of cyclohexane are mixed with 2 moles of hex-1-ene then 

Initial pressure of ${\mathrm{AB}}_3\left(\mathrm{g}\right)$  is 800 mm and total pressure at equilibrium is 900 mm of Hg. Calculate the percentage of ${\mathrm{A}\mathrm{B}}_3$ dissociated in the following reaction:

$\mathrm{A}{\mathrm{B}}_3\left(\mathrm{g}\right)\rightleftharpoons \mathrm{A}{\mathrm{B}}_2\left(\mathrm{g}\right)+\frac{1}{2}{\mathrm{B}}_2\left(\mathrm{g}\right)$