What is the value of ${\mathrm{pK}}_{\mathrm{b}}<mspace\; width="1em"></mspace>({\mathrm{CH}}_3{\mathrm{COO}}^{-}<mspace\; width="1em"></mspace>)$ if ${\mathrm{\lambda }}_{\mathrm{m}}^0=<mspace\; width="1em"></mspace>390<mspace\; width="1em"></mspace>\&<mspace\; width="1em"></mspace>{\mathrm{\lambda }}_{\mathrm{m}}<mspace\; width="1em"></mspace>=<mspace\; width="1em"></mspace>7.8$ for 0.04 of a $<mspace\; width="1em"></mspace>{\mathrm{CH}}_3\mathrm{COOH}$  at $25°\mathrm{C}$ ?

(A) 9.3                                               

(B) 9.2

(C) 4.7                                               

(D) 4.8

Calculate the equilibrium constant (K) for a reaction involving a two-electron transfer,
 if the standard emf of the cell at 25°C is 0.295 V.

$1.<mspace\; width="1em"></mspace>\mathrm{2.95\; \times \; 10⁻²}$

$2.<mspace\; width="1em"></mspace>\mathrm{1.0\; \times \; 10¹}$

3.    2.95 × 10¹⁰

4.   1.0 × 10¹⁰

How much current is necessary to produce $H_2$ gas at the rate of 1 $c{m}^3$ per second under STP ?

(A) 4.305 amp                                           

(B) 17.22 amp

(C) 8.61 amp                                            

(D) 2.1525 amp.

During discharge of a lead storage cell, the density of sulphuric acid in the cell-

1. Increases.                                     

2. Decreases.

3. Remains unchanged.                      

4. Initially increases but decreases subsequently.

-
Why can F₂ gas not be obtained by the electrolysis of any F⁻ salt?

(1) Fluorine is the strongest reducing agent

(2) Fluorine is the strongest oxidising agent.

(3) Fluorine readily reacts with atmospheric O₂

(4) All of the above

The metal that cannot be produced on reduction of its oxide by aluminium is :

(1) K                                                        

(2) Mn

(3) Cr                                                      

(4) Fe

The calomel electrode is reversible with respect to-

1. Mercury                                       

2. $H^{+}$

3. $H{g}^{2+}$                                          

4. $C{l}^{-}$

The standard oxidation potentials are:

Cu → Cu²⁺ + 2e⁻ ; E°ₒₓ = −0.34 V
Cu⁺ → Cu²⁺ + e⁻ ; E°ₒₓ = −0.16 V

Calculate the standard electrode potential for:

Cu⁺ + e⁻ → Cu

1. 0.18V  

2. 0.52V

3. 0.82V    

4. 0.49V

 ${\mathrm{E}}^0<mspace\; width="1em"></mspace>\mathrm{for}<mspace\; width="1em"></mspace>\mathrm{the}<mspace\; width="1em"></mspace>\mathrm{reaction}<mspace\; width="1em"></mspace>$
$\mathrm{Fe}<mspace\; width="1em"></mspace>+<mspace\; width="1em"></mspace>{\mathrm{Zn}}^{2+}<mspace\; width="1em"></mspace>=<mspace\; width="1em"></mspace>\mathrm{Zn}<mspace\; width="1em"></mspace>+<mspace\; width="1em"></mspace>{\mathrm{Fe}}^{2+}<mspace\; width="1em"></mspace>\mathrm{is}<mspace\; width="1em"></mspace>–<mspace\; width="1em"></mspace>0.35<mspace\; width="1em"></mspace>\mathrm{V}.$
The   given   cell   reaction   is: 

1. Feasible  

2. Not feasible

3. In equilibrium 

4. None of the above

The  number of electrons involved in the deposition of 63.5 g of Cu from a solution of $CuS{O}_4$ is:

1 $6.022$ $\times$ ${10}^{23}$ $$

2 $3.011$ $\times$ ${10}^{23}$ $$

3 $12.044$ $\times$ ${10}^{23}$ $$

4 $$ $6.022$ $\times$ ${10}^{22}$