An incorrect statement about equilibrium among the following is:

1.	In equilibrium mixture of ice and water kept in perfectly insulated flask, the mass of ice and water does not change with time.
2.	The intensity of the red colour increases when oxalic acid is added to a solution containing iron (III) nitrate and potassium thiocyanate.
3.	On the addition of a catalyst, the equilibrium constant value is not affected.
4.	The equilibrium constant for a reaction with a negative  ∆H  value decreases as the temperature increases.

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?

At 500 K, equilibrium constant, K_c, for the following reaction is 5.

$\frac{1}{2}{\mathrm{H}}_2\left(\mathrm{g}\right)+\frac{1}{2}{\mathrm{I}}_2\left(\mathrm{g}\right)\rightleftharpoons \mathrm{HI}\left(\mathrm{g}\right)$

What would be the equilibrium constant K_c for the reaction?

$2\mathrm{HI}\left(\mathrm{g}\right)\rightleftharpoons {\mathrm{H}}_2\left(\mathrm{g}\right)+{\mathrm{I}}_2\left(\mathrm{g}\right)$

1.  0.04

2.  0.4

3.  25

4.  2.5

 At 450 K, K_p= 2.0 × 10¹⁰ bar^-1 for the given reaction at equilibrium

$2{{\mathrm{SO}}_2}_{(\mathrm{g})\hspace{0.17em}}+$ ${{\mathrm{O}}_2}_{\left(\mathrm{g}\right)}$ $\rightleftharpoons$ $2{{\mathrm{SO}}_3}_{\left(\mathrm{g}\right)}$ $$

The value of  K_c at this temperature would be :

$1.$ $7.48$ $\times {10}^{12}$ ${\mathrm{M}}^{-1}$
$2.$ $6.56$ $\times$ ${10}^{11}$ ${\mathrm{M}}^{-1}$
$3.$ $7.48$ $\times$ ${10}^{11}$ ${\mathrm{M}}^{-1}$
$4.$ $1.23$ $\times$ ${10}^{10}$ ${\mathrm{M}}^{-1}$