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:
\(\mathrm{MgCO}_{3(\mathrm{~s})} \rightleftharpoons \mathrm{MgO}_{(\mathrm{s})}+\mathrm{CO}_{2(\mathrm{~g})}\)
1. \(\mathrm{K_P} =\mathrm{P}_{\mathrm{CO}_2}\)
2. \(\mathrm{K_P} =\mathrm{P}_{\mathrm{CO}_2} \times \frac{\mathrm{P}_{\mathrm{CO}_2 \times \mathrm{P}_{\mathrm{MgO}^{\mathrm{}}}}}{\mathrm{P}_{\mathrm{MgCO}_3}}\)
3. \(\mathrm{K_P} =\frac{\mathrm{P}_{\mathrm{CO}_2}+\mathrm{P}_{\mathrm{MgO}^{\mathrm{}}}}{\mathrm{P}_{\mathrm{MgCO}_3}}\)
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. $K{a}_1=\frac{1}{K{a}_2}$

What happens to the equilibrium constant of a reversible reaction
if the concentration of reactants is increased?

1. It depends on the amount of concentration
2. It remains unchanged
3. It decrease
4. It increase

Conjugate acid of NH₂^– is:

1. NH₄OH

2. NH₄⁺

3. \(NH_{2}^{-}\)

4. NH₃

Incorrect statement about pH and H⁺ is: 

Given the following two reactions:
     A + B ⇌ C + D, with rate constant K₁
     E + F ⇌ G + H, with rate constant K₂

If C + D + E + F produces a product, what is the rate constant for this reaction?