The equilibrium pressure of C₂H₆ when it is placed in a flask at 4.0 atm pressure at  899 K would be:

C₂H_6 (g)  $\leftrightharpoons$ C₂H_4 (g) + H_2 (g)

(K_p = 0.04 atm at 899 K)

1.	4.12 atm	2.	3.62 atm
3.	1.54 atm	4.	2.16 atm

A sample of pure PCl₅ was introduced into an evacuated vessel at 473 K.
After equilibrium was attained, a concentration of PCl₅ 
was found to be 0.5 × 10^–1 mol L^–1. If the value of
 K_c is 8.3 × 10^–3 mol L^–1, the concentrations of
PCl₃ and Cl₂ at equilibrium would be:

PCl_5 (g)  $\leftrightharpoons \hspace{0.17em}$PCl_3 (g) + Cl_2(g)
 

An incorrect statement about equilibrium among the following is:

When hydrochloric acid is added to cobalt nitrate solution at room temperature, the following reaction takes place and the reaction mixture becomes blue. On cooling the mixture it becomes pink
\([Co(H_2O)_6]^{3+} (aq) + 4Cl^-(aq) \rightleftharpoons [CoCl_4]^{2-} (aq) + 6H_2O(l)\\ ~~~\small{(Pink})~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~\small{ (Blue)}\)

On the basis of the information given above, mark the correct answer:

The pH of neutral water at $25°\mathrm{C}$ is 7.0. As the temperature increases, ionisation of water increases. However, the concentration of H⁺ ions and OH^- ions is equal. What will be the pH of pure water at 60°C ?

${\mathrm{K}}_{{\mathrm{a}}_1},$ ${\mathrm{K}}_{{\mathrm{a}}_2}$ $\mathrm{and}$ ${\mathrm{K}}_{{\mathrm{a}}_3}$ are the respective ionisation constants for the following reactions.
\(\mathrm{H}_2 \mathrm{~S} \rightleftharpoons \mathrm{H}^{+}+\mathrm{HS}^{-}\)
\(\mathrm{HS}^{-} \rightleftharpoons \mathrm{H}^{+}+\mathrm{S}^{2-}\)
\(\mathrm{H}_2 \mathrm{~S} \rightleftharpoons 2 \mathrm{H}^{+}+\mathrm{S}^{2-}\)
The correct relationship between ${\mathrm{K}}_{{\mathrm{a}}_1},$ ${\mathrm{K}}_{{\mathrm{a}}_2}$ $\mathrm{and}$ ${\mathrm{K}}_{{\mathrm{a}}_3}$ is:
1. \(\mathrm{K}_{\mathrm{a}_3}=\mathrm{K}_{\mathrm{a}_1} \times \mathrm{K}_{\mathrm{a}_2} \)
2. \(\mathrm{K}_{\mathrm{a}_3}=\mathrm{K}_{\mathrm{a}_1}+\mathrm{K}_{\mathrm{a}_2} \)
3. \(K_{a_3}=K_{a_1}-K_{a_2} \)
4. \(\mathrm{K}_{\mathrm{a}_3}=\mathrm{K}_{\mathrm{a}_1} / \mathrm{K}_{\mathrm{a}_2}\)

The mixture that will produce a buffer solution when mixed in equal volumes is:

${\mathrm{K}}_{\mathrm{a}}$ for ${\mathrm{CH}}_3\mathrm{COOH}$ is $1.8\times {10}^{-5}$ and K_b for ${\mathrm{NH}}_4\mathrm{OH}$ is $1.8\times {10}^{-5}$. The pH of ammonium acetate will be: 

1. 7.005

2. 4.75

3. 7.0

4. Between 6 and 7

On increasing the pressure, the direction in which the gas phase reaction proceeds to re-establish equilibrium is predicted by applying Le-Chatelier's principle. Consider the reaction,

${\mathrm{N}}_2\left(\mathrm{g}\right)+3{\mathrm{H}}_2\left(\mathrm{g}\right)\rightleftharpoons 2{\mathrm{NH}}_3\left(\mathrm{g}\right)$

Which of the following is correct, if the total pressure at which the equilibrium is established is increased without changing the temperature?