An aqueous solution of hydrogen sulphide shows the equilibrium,

           H₂S $\rightleftharpoons$H⁺ + HS^-

If dilute hydrochloric acid is added to an aqueous solution of hydrogen sulphide without any temperature change, then:

1.	The equilibrium constant will change
2.	The concentration of HS- will increase
3.	The concentration of undissociated hydrogen sulphide will decrease
4.	The concentration of HS- will decrease

When HCl gas is passed through a saturated solution of common salt, pure NaCl is precipitated because:

1. the impurities dissolve in HCl

2. HCl is highly soluble in H₂O

3. the product of [Na⁺] and [Cl^-] exceeds the solubility product of NaCl

4. the solubility product of NaCl is lowered by the chloride ion from aqueous HCl

Which combination of acid and base forms a salt that gives an alkaline aqueous solution?

1. A strong acid and a strong base

2. A strong acid and a weak base

3. A weak acid and a weak base

4. A weak acid and a strong base

40% of a racemic mixture of 0.2 moles of N₂ and 0.6 moles of H₂ react to give NH₃ according to the equation, N₂ (g) + H₂(g)$\rightleftharpoons$2NH₃ (g) at constant temperature and pressure. Then the ratio of the final volume to the initial volume of gases is:

1. 4:5                       

2. 5:4

3. 7:10                     

4. 8:5

On addition of inert gas at constant volume to the reaction, N₂ + 3H₂$\rightleftharpoons$2NH₃ at  equilibrium:

1. the reaction halts

2. forward reaction is favored

3. the reaction remains unaffected

4. backward reaction is favored

In which of the following case reaction goes farthest to completion?

1. K=10³                                         

2. K=10^-2

3. K=10                                           

4. K=10⁰

For the reversible reaction,

N₂ (g) + 3H₂(g) $\rightleftharpoons$2NH₃ (g) + heat

According to Le Chatelier’s principle, the equilibrium will shift in the forward direction by:                   

1. Increasing the concentration of NH₃ (g) 

2. Decreasing the pressure

3. Decreasing the concentrations of N₂ (g)and H₂(g)

4. Increasing pressure and decreasing temperature

The equilibrium constants for the reactions are:

H₃PO₄ $\stackrel{K_{\mathit{1}}}{\rightleftharpoons }$H+ + H₂P${\mathrm{O}}_4^{-}$;                      K₁

H₂P${\mathrm{O}}_4^{-}$$\stackrel{K_{\mathit{2}}}{\rightleftharpoons }$H+ + HP${\mathrm{O}}_4^{2-}$;                      K₂

HP${\mathrm{O}}_4^{2-}$$\stackrel{K_{\mathit{3}}}{\rightleftharpoons }$H+ + P${\mathrm{O}}_4^{3-}$                        K₃

The equilibrium constants for 

H₃PO₄$\rightleftharpoons$3H+ + P${\mathrm{O}}_4^{3-}$  will be:

1. K₁/K₂K₃                                         

2. K₁xK₂xK₃

3. K₂/K₁K₃                                       

4. K₁+K₂+K₃

The equilibrium constant for the reaction,

SO₃ (g) $\rightleftharpoons$SO₂(g) + $\frac{1}{2}$O₂(g); K_c = 4.9 x 10^-2.

The K_c for the reaction: 2SO₂(g) + O₂(g) $\rightleftharpoons$2SO₃(g) will be:

1. 416.5                             

2. 2.40 x10^-3

3. 9.8 x10^-2                       

4. 4.9 x10^-2

The correct representation for the solubility product constant of Ag₂CrO₄ is:

1. [Ag⁺]²[${\mathrm{CrO}}_4^{2-}$]                             

2. [Ag²⁺][${\mathrm{CrO}}_4^{2-}$]

3. [2Ag⁺][${\mathrm{CrO}}_4^{2-}$]                             

4. [2Ag⁺]²[${\mathrm{CrO}}_4^{2-}$]