10 gm CaCO₃ is taken in a one litre container. The active mass of CaCO₃ is: 

1. 0.1

2. 1

3. zero

4. 10

The most hydrolyzed salt among the following is-

(Assume that K_b of all weak bases is the same)
1. NH₄Cl
2. CuSO₄
3. AlCl₃
4. All are equally hydrolyzed.

The increasing order of basic strength of Cl^-, ${\mathrm{CO}}_3^{2-}<mspace\; width="1em"></mspace>$, CH₃COO^-, OH^-, F^- is:

1. Cl^-<F^-<CH₃COO^-<${\mathrm{CO}}_3^{2-}<mspace\; width="1em"></mspace>$<OH^-

2.  Cl^-<F^-< ${\mathrm{CO}}_3^{2-}<mspace\; width="1em"></mspace>$<CH₃COO^-<OH^-

3. CH₃COO^-<Cl^-<F^-<${\mathrm{CO}}_3^{2-}<mspace\; width="1em"></mspace>$<OH^-

4. none of the above

KMnO₄ can be prepared from K₂MnO₄ as per reaction,

 $3{\mathrm{MnO}}_4^{2-}<mspace\; width="1em"></mspace>+<mspace\; width="1em"></mspace>2{\mathrm{H}}_2\mathrm{O}<mspace\; width="1em"></mspace>\rightleftharpoons <mspace\; width="1em"></mspace>2{\mathrm{MnO}}_4^{-}<mspace\; width="1em"></mspace>+<mspace\; width="1em"></mspace>{\mathrm{MnO}}_2<mspace\; width="1em"></mspace>+<mspace\; width="1em"></mspace>4{\mathrm{OH}}^{-}$

The reaction can go to completion by removing OH^- ions by adding             

1. HCl

2. KOH

3. CO₂

4. SO₂

The equilibrium constant of the reaction A₂(g) + B₂(g) $\rightleftharpoons$ 2AB (g) at 100ºC is 50.  If a one litre flask containing one mole of A₂ is connected to a two litre flask containing two moles of B₂, how many moles of AB will be formed at 373 K ?

1. 2.8

2. 1.9

3. 2.1

4. 3.6

The solubility of AgCl in 0.2 M magnesium chloride will be $\left({\mathrm{K}}_{sp}<mspace\; width="1em"></mspace>\mathrm{of}<mspace\; width="1em"></mspace>\mathrm{AgCl}<mspace\; width="1em"></mspace>=<mspace\; width="1em"></mspace>1.8<mspace\; width="1em"></mspace>\mathrm{x}<mspace\; width="1em"></mspace>{10}^{-10}\right)$

1.  1.8 x ${10}^{-10}$

2.  1.8 x ${10}^{-11}$

3.  9 x ${10}^{-10}$

4.  4.5 x ${10}^{-10}$

Which of the following compounds is in the correct sequence in terms of relative basic strength?

1. ${\mathrm{C}}_2{\mathrm{H}}_5{\mathrm{O}}^{-}>\mathrm{CH}\equiv {\mathrm{C}}^{-}>{\mathrm{OH}}^{-}$

2. $\mathrm{CH}\equiv {\mathrm{C}}^{-}{>}^{-}\mathrm{OH}>{\mathrm{C}}_2{\mathrm{H}}_5{\mathrm{O}}^{-}$

3. $\mathrm{CH}\equiv {\mathrm{C}}^{-}>{\mathrm{C}}_2{\mathrm{H}}_5{\mathrm{O}}^{-}>{\mathrm{OH}}^{-}$

4. ${\mathrm{C}}_2{\mathrm{H}}_5{\mathrm{O}}^{-}>{\mathrm{OH}}^{-}>\mathrm{CH}\equiv {\mathrm{C}}^{-}$

Among the following examples, the species that behave(s) as a Lewis acid is/are: 
\(\mathrm{BF}_3, \mathrm{SnCl}_2, \mathrm{SnCl}_4\)

1. Stannous chloride, Stannic chloride
2. $B{F}_3$, Stannous chloride
3. Only $B{F}_3$
4. $B{F}_3$, Stannous chloride, Stannic chloride

The rapid change of pH near the stoichiometric point of an acid-base titration is the basis of indicator detection. pH of the solution is related to the ratio of the concentrations of the conjugate acid (HIn) and base (In–) forms of the indicator by the expression:

1. $\log \frac{\left[\mathrm{HIn}\right]}{\left[{\mathrm{In}}^{-}\right]}={\mathrm{pK}}_{\mathrm{In}}-\mathrm{pH}$

2. $\log \frac{\left[\mathrm{HIn}\right]}{\left[{\mathrm{In}}^{-}\right]}=\mathrm{pH}-{\mathrm{pK}}_{\mathrm{In}}$

3. $\log \frac{\left[{\mathrm{In}}^{-}\right]}{\left[\mathrm{HIn}\right]}=-pH\; +\hspace{0.17em}{\mathrm{pK}}_{\mathrm{In }}$

4. All of the above.

Case Study:

A buffer solution is defined as a solution whose pH remains practically constant even when small amounts of an acid or a base are added to it. 
Types of buffer solutions : 
(i) Acidic buffer: It is a solution of a mixture of a weak acid and a salt of this weak acid with a strong base (e.g. CH₃COOH + CH₃COONa)
(ii) Basic buffer: It is the solution of a mixture of a weak base and a salt of this weak base with a strong acid (e.g. NH₄OH + NH₄Cl)
Henderson's equation is used to determine pH of buffer mixtures of differ types: 
for acidic buffer Henderson's equation is 
pH= pK_a + log \([Salt] \over [Acid]\)   (k_a = ionisation constant of weak acid )

for basic buffer Henderson's equation is : 

POH = Pk_b + log \([Salt] \over [Base]\)   (k_b = ionisation constant of weak base )

How many moles of HCl are required with 0.01 mole NaCN to prepare a buffer solution of pH =9? 
(k_a of HCN = \(1 \times 10^{-10}\))

1. 0.009 

2. 0.09 

3. 0.9 

4. Buffer solution cannot formed