One mole of H₂O and one mole of CO are taken in a 10 L vessel and heated to
725 K. At equilibrium, 40% of water (by mass) reacts with CO according to the equation,

H₂O _(g) + CO _(g)  $\leftrightharpoons$H_2 (g) + CO_2 (g)

The equilibrium constant for the above-mentioned reaction would be:

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)

PCl_5 (g)  $\leftrightharpoons$PCl_3 (g) + Cl_2(g)

 $$\begin{gathered} 1. \left[{\mathrm{PCl}}_3\right] = 0.02 \mathrm{mol} {\mathrm{L}}^{-1} , \left[{\mathrm{Cl}}_2\right] = 0.04 \mathrm{mol} {\mathrm{L}}^{-1} \\ 2. \left[{\mathrm{PCl}}_3\right]=\left[{\mathrm{Cl}}_2\right] = 0.02 \mathrm{mol} {\mathrm{L}}^{-1} \\ 3. \left[{\mathrm{PCl}}_3\right] = 0.04 \mathrm{mol} {\mathrm{L}}^{-1}, \left[{\mathrm{Cl}}_2\right] =0.02 \mathrm{mol} {\mathrm{L}}^{-1} \\ 4. \left[{\mathrm{PCl}}_3\right]= \left[{\mathrm{Cl}}_2\right] = 0.04 \mathrm{mol} {\mathrm{L}}^{-1} \end{gathered}$$

An incorrect statement about equilibrium among the following is:

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. On the basis of this information mark the correct answer.

$$\begin{gathered} \underset{\left(\mathrm{Pink}\right)}{{\left[\mathrm{Co}\right({\mathrm{H}}_2{\mathrm{O}}_6\left)\right]}^{3+}}\left(\mathrm{aq}\right)+4{\mathrm{Cl}}^{-}\left(\mathrm{aq}\right)\rightleftharpoons \\ \underset{\left(\mathrm{Blue}\right)}{{\left[{\mathrm{CoCl}}_4\right]}^{2-}}\left(\mathrm{aq}\right)+6{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right) \end{gathered}$$

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. The pH of pure water at $60°\mathrm{C}$ is:

${\mathrm{K}}_{{\mathrm{a}}_1}, {\mathrm{K}}_{{\mathrm{a}}_2} \mathrm{and} {\mathrm{K}}_{{\mathrm{a}}_3}$ are the respective ionisation constants for the following reactions.

$$\begin{gathered} {\mathrm{H}}_2\mathrm{S}\rightleftharpoons {\mathrm{H}}^{+}+{\mathrm{HS}}^{-} \\ {\mathrm{HS}}^{-}\rightleftharpoons {\mathrm{H}}^{+}+{\mathrm{S}}^{2-} \\ {\mathrm{H}}_2\mathrm{S}\rightleftharpoons 2{\mathrm{H}}^{+}+{\mathrm{S}}^{2-} \end{gathered}$$

The correct relationship between ${\mathrm{K}}_{{\mathrm{a}}_1}, {\mathrm{K}}_{{\mathrm{a}}_2} \mathrm{and} {\mathrm{K}}_{{\mathrm{a}}_3}$ is

Which of the following mixtures will produce a buffer solution when mixed in equal volumes?

${\mathrm{K}}_{\mathrm{a}}$ 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 

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

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

If the total pressure at which the equilibrium is established is increased without changing the temperature, then the correct option is:

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?

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

The addition of a small amount of argon at constant volume will not affect the equilibrium among the following:

 At 450 K, K_p= 2.0 × 10¹⁰ bar^-1 for the given reaction at equilibrium $2{{\mathrm{SO}}_2}_{(\mathrm{g}) }+ {{\mathrm{O}}_2}_{\left(\mathrm{g}\right) }\rightleftharpoons 2{{\mathrm{SO}}_3}_{\left(\mathrm{g}\right) }$

The value of  K_c at this temperature would be-

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

For the reaction: ${\mathrm{FeO}}_{\left(\mathrm{s}\right)} + {\mathrm{CO}}_{\left(\mathrm{g}\right)} \rightleftharpoons {\mathrm{Fe}}_{\left(\mathrm{s}\right)} \hspace{0.17em}+\hspace{0.17em}{{\mathrm{CO}}_2}_{\left(\mathrm{g}\right)}, {\mathrm{K}}_{\mathrm{p}} = 0.265$  at 1050 K. If the initial partial pressures are p_CO= 1.4 atm and ${\mathrm{p}}_{{\mathrm{CO}}_2}$= 0.80 atm, the partial pressure of CO₂ at equilibrium at 1050 K would be:

$$\begin{gathered} 1. 4.61 \mathrm{atm} \\ 2. 1.74 \mathrm{atm} \\ 3. 0.46 \mathrm{atm} \\ 4. 0.17 \mathrm{atm} \end{gathered}$$

For the reaction, 

$$\begin{gathered} {\mathrm{N}}_2 \left(\mathrm{g}\right) + 3{\mathrm{H}}_2\left(\mathrm{g}\right) \rightleftharpoons 2{\mathrm{NH}}_3 \left(\mathrm{g}\right); \\ {\mathrm{K}}_{\mathrm{c}}= 0.061 {\mathrm{mol}}_{}^{-2}{\mathrm{L}}^2 \end{gathered}$$  at 500K. At a particular instant of time, [N₂] = 3.0 mol L^–1, [H₂] = 2.0 mol L^–1  and [NH₃] = 0.5 mol L^–1 .

Which of the following is true?

At 1127 K and 1 atm pressure, a gaseous mixture of CO and CO₂ in equilibrium with solid carbon has 90.55% CO by mass. 

${\mathrm{C}}_{\left(\mathrm{s}\right)} + {{\mathrm{CO}}_2}_{\left(\mathrm{g}\right)} \rightleftharpoons 2{\mathrm{CO}}_{\left(\mathrm{g}\right)}$

At the specified temperature, K_c for this reaction would be

$$\begin{gathered} 1. 0.25 \mathrm{mol} {\mathrm{L}}^{-1} \\ 2. 0.34 \mathrm{mol} {\mathrm{L}}^{-1} \\ 3. 0.15 \mathrm{mol} {\mathrm{L}}^{-1} \\ 4. 1.25 \mathrm{mol} {\mathrm{L}}^{-1} \end{gathered}$$

Consider the following reaction.

If a reaction vessel at 400°C is filled with an equimolar mixture of CO and steam such that ${\mathrm{p}}_{\mathrm{CO}} = {\mathrm{p}}_{{\mathrm{H}}_2\mathrm{O}} = 4.0 \mathrm{bar}$ , the partial pressure of H₂ at equilibrium is:  

${\mathrm{CO}}_{\left(\mathrm{g}\right)} + {\mathrm{H}}_2{\mathrm{O}}_{\left(\mathrm{g}\right)} \rightleftharpoons {{\mathrm{CO}}_2}_{\left(\mathrm{g}\right)}+ {{\mathrm{H}}_2}_{\left(\mathrm{g}\right)}$ ;   K_p= 10.1 at 400°C

$$\begin{gathered} 1. 30.40 \mathrm{bar} \\ 2. 3.04 \mathrm{bar} \\ 3. 1.24 \mathrm{bar} \\ 4. 3.04 \mathrm{atm} \end{gathered}$$

$$\begin{gathered} \left(\mathrm{a}\right)\hspace{0.17em}{{\mathrm{Cl}}_2}_{\left(\mathrm{g}\right)} \rightleftharpoons 2{\mathrm{Cl}}_{\left(\mathrm{g}\right)} ; \\ {\mathrm{K}}_{\mathrm{c} }= 5.0 \times {10}^{-39} \mathrm{mol}/\mathrm{L} \\ \left(\mathrm{b}\right) {{\mathrm{Cl}}_2}_{\left(\mathrm{g}\right)} + 2{\mathrm{NO}}_{\left(\mathrm{g}\right)} \rightleftharpoons 2{\mathrm{NOCl}}_{\left(\mathrm{g}\right)} ; \\ {\mathrm{K}}_{\mathrm{c}} = 3.7 \times {10}^8 \mathrm{mol}/\mathrm{L} \\ \left(\mathrm{c}\right) {{\mathrm{Cl}}_2}_{\left(\mathrm{g}\right)}+ 2{{\mathrm{NO}}_2}_{\left(\mathrm{g}\right)} \rightleftharpoons 2{\mathrm{NO}}_2{\mathrm{Cl}}_{ \left(\mathrm{g}\right)} ; \\ {\mathrm{K}}_{\mathrm{c}} = 1.8 \mathrm{mol}/\mathrm{L} \end{gathered}$$

The reaction(s) in which there will be an appreciable concentration of reactants and products, among the following reaction is 

For the reaction, $3{\mathrm{O}}_2 \left(\mathrm{g}\right) \rightleftharpoons 2{\mathrm{O}}_3 \left(\mathrm{g}\right), {\mathrm{K}}_{\mathrm{c}} = 2.0 \times {10}^{-50}$ L/mol at 25^°C. If the equilibrium concentration of O₂ in air at 25^°C is $1.6 \times {10}^{-2}$, the concentration of O₃ would be 

$$\begin{gathered} 1. 2.86 \times {10}^{28} \mathrm{M} \\ 2. 28.66 \times {10}^{-28 }\mathrm{M} \\ 3. 1.43 \times {10}^{-28} \mathrm{M} \\ 4. 2.86 \times {10}^{-28} \mathrm{M} \end{gathered}$$

Consider the following reaction:

If the reaction mixture contains 0.30 mol of CO, 0.10 mol of H_2, and 0.02 mol of H₂O, and an unknown amount of CH₄ at equilibrium at 1300 K in a 1L flask, the concentration of CH₄ in the mixture will be -

${\mathrm{CO}}_{\left(\mathrm{g}\right)} + 3{{\mathrm{H}}_2}_{\left(\mathrm{g}\right) }\rightleftharpoons {{\mathrm{CH}}_4}_{\left(\mathrm{g}\right) } + {\mathrm{H}}_2{\mathrm{O}}_{\left(\mathrm{g}\right) }$

(K_c = 3.90 at 1300 K) 

$$\begin{gathered} 1. 5.85 \times {10}^2 \mathrm{M} \\ 2. 5.85 \times {10}^{-1} \mathrm{M} \\ 3. 5.85 \times {10}^3 \mathrm{M} \\ 4. 5.85 \times {10}^{-2} \mathrm{M} \end{gathered}$$

The concentration of hydrogen ion in a sample of soft drink is $3.8 \times {10}^{-3} \mathrm{M}$. The  pH of a soft drink will be:

The pH of a sample of vinegar is 3.76. The concentration of H⁺ ions in the sample will be -

$$\begin{gathered} 1. 1.74 \times {10}^{-4 }\mathrm{M} \\ 2. 2.34 \times {10}^{-11}\mathrm{M} \\ 3. 2.86 \times {10}^{-10} \mathrm{M} \\ 4. 1.74 \times {10}^{-3} \mathrm{M} \end{gathered}$$

$$\begin{gathered} 1. 8.95 \\ 2. 1.43 \\ 3. 3.56 \\ 4. 6.30 \end{gathered}$$

A solution containing 2 g of TlOH in 2 litres of water has a pH of?

(Atomic wt of Tl =204)

When 1mL of 13.6 M HCl is diluted with water to give 1 litre of solution, the pH of the resultant solution will be:

The degree of ionization of a 0.1M bromoacetic acid solution is 0.132. The pK_a of bromoacetic acid solution will be

$$\begin{gathered} 1. 1.65 \\ 2. 2.75 \\ 3. 11.25 \\ 4. 2.57 \end{gathered}$$

A 0.001 M aniline solution has a pH of?

( ${\mathrm{K}}_{\mathrm{b}} = 4.27 \times {10}^{-10}$ )

$$\begin{gathered} 1. 6.19 \\ 2. 7.81 \\ 3. 8.34 \\ 4. 9.81 \end{gathered}$$

The percentage of 0.02 M dimethylamine solution ionized if solution also contains 0.1 M NaOH solution (K_b of dimethylamine = 5.4 × 10^–4 ) will be -

0.561 g of KOH is dissolved in water to give 200 mL of solution at 298 K. The pH of the solution will be

At 298 K, the solubility of Sr(OH)₂ is 19.23 g/L. The pH of the solution will be:

The pH of a 0.1 M solution of cyanic acid (HCNO) is 2.34. The ionization constant of the acid will be:

$$\begin{gathered} 1. 2.02 \times {10}^4 \\ 2. 3.14 \times {10}^3 \\ 3. 2.02 \times {10}^{-4} \\ 4. 1.01 \times {10}^{-4} \end{gathered}$$

The ionization constant of nitrous acid is 4.5 × 10^–4. The pH of a 0.04 M sodium nitrite solution will be:

The ionization constant of chloroacetic acid is 1.35 × 10^–3. The pH of a 0.1 M acid solution will be:

The ionic product of water at 310 K is 2.7 × 10^–14.

The pH of neutral water at this temperature will be:

A mixture of 10 mL of 0.2 M Ca(OH)2 and 25 mL of 0.1 M HCl is prepared. The pH of the resultant mixture would be:

(For iron sulphide, K_sp = 6.3 × 10^–18). 

The maximum concentration of equimolar solutions of ferrous sulphate and sodium sulphide so that when mixed in equal volumes, there is no precipitation of iron sulphide will be-

$$\begin{gathered} 1. 5.02 \times {10}^{-9} \mathrm{M} \\ 2. 5.02 \times {10}^{9 }\mathrm{M} \\ 3. 2. 25 \times {10}^{-13} \mathrm{M} \\ 4. \mathrm{Can}\text{'}\mathrm{t} \mathrm{predict} \end{gathered}$$

Which of the following salt solutions is basic in nature?

Given that the equilibrium constant for the reaction 

$2S{O}_{2\left(g\right)} + O_{2\left(g\right)} \leftrightharpoons 2S{O}_{3\left(g\right) }$

has a value of 278 at a particular temperature. The value of the equilibrium constant for the following reaction at the same temperature is: 

$S{O}_{3\left(g\right)} \leftrightharpoons S{O}_{2\left(g\right) }+ 1/2 O_{2\left(g\right) }$

Given reaction: $A_{2\left(g\right)} + B_{2\left(g\right) }\leftrightharpoons 2A{B}_{\left(g\right) }$

At equilibrium , the concentrations of ${\mathrm{A}}_2 = 3.0 \times {10}^{-3} \mathrm{M} ; {\mathrm{B}}_2 = 4.2 \times {10}^{-3} \mathrm{M} \mathrm{and} \mathrm{AB} = 2.8 \times {10}^{-3} \mathrm{M}$

If the reaction takes place in a sealed vessel at $527°C$, then the value of $K_C$ will be : 

In qualitative analysis, the metals of Group I can be separated from other ions by precipitating them as chloride salts. A solution initially contains Ag⁺
and Pb²⁺ at a concentration of 0.10 M. Aqueous HCl is added to this solution until the Cl– concentration is 0.10 M. What will the concentration of Ag⁺
and Pb²⁺ be at equilibrium?

$$\begin{gathered} \left({\mathrm{K}}_{\mathrm{sp}} \mathrm{for} \mathrm{AgCl}=1.8\times {10}^{-10}\right) \\ (K_{sp} for PbC{l}_2=1.7\times {10}^{-5}) \end{gathered}$$

    $\left[P{b}^{2+}\right]=1.7\times {10}^{-4}M$

    $\left[P{b}^{2+}\right]=1.7\times {10}^{-6}M$

   $\left[P{b}^{2+}\right]=8.5\times {10}^{-5}M$

   $\left[P{b}^{2+}\right]=1.7\times {10}^{-3}M$

The reaction

$2A+B\left(g\right) \rightleftharpoons 3C\left(g\right)+D\left(g\right)$

is begun with the concentrations of A and B both at an intial value of 1.00 M. When equilibrium is reached, the concentration of D is measured and found to be 0.25 M. The value for the equilibrium constant for this reaction is given by the expression.

The equilibrium constant Kp for the following reaction is-

${{\mathrm{MgCO}}_3}_{\left(\mathrm{s}\right)}\rightleftharpoons {\mathrm{MgO}}_{\left(\mathrm{s}\right)}+{{\mathrm{CO}}_2}_{\left(\mathrm{g}\right)}$

Correct relation between dissociation constant's of a di-basic acid :

For any reversible reaction, if we increase the concentration of  the reactants, then effect on equilibrium constant :

1. Depends on the amount of concentration
2. Unchanged
3. Decrease
4. Increase

Conjugate acid of NH₂^–:

2. NH₄⁺

Incorrect statement about pH and H⁺ is : 

$$\begin{gathered} A + B \leftrightharpoons C + D Cons\tan t = K_1 \\ E + F \rightleftharpoons G + H Cons\tan t = K_2 \end{gathered}$$

then C + D + E + F ⇒ product. The constant of reaction will be :

The fertilizer which makes the soil acidic :

Among the following examples , the species that behave as a lewis acid  are : 
$B{F}_3 , SnC{l}_2 , SnC{l}_4$