Consider the following reaction:
A₂(g) + B₂(g)  ⇋ 2AB(g)  
At equilibrium, the concentrations of  A₂ = 3.0×10^–3 M;  B₂ = 4.2×10^–3 M and AB = 2.8×10^–3M.
The value \(K_C\)  for the above-given reaction in a sealed container at 527°C is:

1.	3.9	2.	0.6
3.	4.5	4.	2.0

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⁺ at equilibrium?

(K_sp for AgCl  = 1.8 × 10^-10)$$
(K_sp for PbCl₂ = 1.7 × 10^-5) 

The reaction-

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

begins with the concentrations of A and B both at an initial 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:

1. $\left[{(0.75)}^3\right(0.25\left)\right]÷\left[{(0.50)}^2\right(0.75\left)\right]$

2. $\left[{(0.75)}^3\right(0.25\left)\right]÷\left[{(0.50)}^2\right(0.25\left)\right]$

3. $\left[{(0.75)}^3\right(0.25\left)\right]÷\left[{(0.75)}^2\right(0.25\left)\right]$

4. $\left[{(0.75)}^3\right(0.25\left)\right]÷\left[{(1.00)}^2\right(1.00\left)\right]$

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 will be:

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

1.  $3.6$ $\times$ ${10}^{-3}$ $$

2.  $6.0$ $\times$ ${10}^{-2}$ $$

3.  $1.3$ $\times$ ${10}^{-5}$ $$

4.  $1.8$ $\times$ ${10}^{-3}$ $$

The equilibrium constant Kp for the following reaction is:

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

1. $\mathrm{Kp }= {\mathrm{P}}_{{\mathrm{CO}}_2}$

2. $\mathrm{Kp}={\mathrm{P}}_{{\mathrm{CO}}_2}\times \frac{{\mathrm{P}}_{{\mathrm{CO}}_2}\times {\mathrm{P}}_{\mathrm{MgO}}}{{\mathrm{P}}_{{\mathrm{MgCO}}_3}}$

3. $\mathrm{Kp}=\frac{{\mathrm{P}}_{{\mathrm{CO}}_2}+{\mathrm{P}}_{\mathrm{MgO}}}{{\mathrm{P}}_{{\mathrm{MgCO}}_3}}$

4. $\mathrm{Kp}=\frac{{\mathrm{P}}_{{\mathrm{MgCO}}_3}}{{\mathrm{P}}_{{\mathrm{CO}}_2}\times {\mathrm{P}}_{\mathrm{MgO}}}$

The correct relation between dissociation constants of a di-basic acid is:

1. $K{a}_1=K{a}_2$

2. $K{a}_1>K{a}_2$

3. $K{a}_1<K{a}_2$

4. $K{a}_1=\frac{1}{K{a}_2}$

For any reversible reaction, if we increase the concentration of the reactants, the effect on equilibrium constant will:
1. Depend on the amount of concentration
2. Remain unchanged
3. Decrease
4. Increase

Conjugate acid of NH₂^– is:

1. NH₄OH

2. NH₄⁺

3. \(NH_{2}^{-}\)

4. NH₃

Incorrect statement about pH and H⁺ is: 

$\mathrm{If\; A}$ $+$ $B$ $\leftrightharpoons$ $C$ $+$ $D$ $Cons\tan t$ $=$ $K_1$
$E$ $+$ $F$ $\rightleftharpoons$ $G$ $+$ $H$ $Cons\tan t$ $=$ $K_2$

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

1.  $\frac{K_1}{K_2}$

2.  $\frac{K_2}{K1}$

3. $K_1{K}_2$

4. None of these