What is the molarity of the standard solution if the solubility product for a salt of type AB is $4\times {10}^{-8}$?

1. $2\times {10}^{-4}$ $\mathrm{mol}/\mathrm{L}$

2. $16\times {10}^{-16}$ $\mathrm{mol}/\mathrm{L}$

3. $2\times {10}^{-16}$ $\mathrm{mol}/\mathrm{L}$

4. $4\times {10}^{-4}$ $\mathrm{mol}/\mathrm{L}$

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.  $1.8$ $\times$ ${10}^{-3}$ $$

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:

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

Which of the following salt solutions is basic in nature?

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

Equilibrium constants K₁ and K₂ for the following equilibria

$\begin{array}{l}\mathrm{NO}\left(\mathrm{g}\right)+\frac{1}{2}{\mathrm{O}}_2\stackrel{{\mathrm{K}}_1}{⟶}{\mathrm{NO}}_2\left(\mathrm{g}\right)\text{\hspace{0.17em}and\hspace{0.17em}}\\ 2{\mathrm{NO}}_2\left(\mathrm{g}\right)\stackrel{{\mathrm{K}}_2}{\rightleftharpoons }2\mathrm{NO}\left(\mathrm{g}\right)+{\mathrm{O}}_2\left(\mathrm{g}\right)\end{array}$

are related as:

1. ${\mathrm{K}}_2=\frac{1}{{\mathrm{K}}_1}$

2. ${\mathrm{K}}_2=\frac{{\mathrm{K}}_1}{2}$

3. ${\mathrm{K}}_2=\frac{1}{{\mathrm{K}}_1^2}$

4. ${\mathrm{K}}_2={\mathrm{K}}_1^2$