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 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.

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 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 are in the correct sequence in terms of relative basic strength?

1. \(C_{2} H_{5} O^{-} > CH \equiv C^{-} > \left(OH\right)^{-}\)

2. \(CH \equiv C^{-} >(OH)^{-} > C_{2} H_{5} O^{-}\)

3. \(CH \equiv C^{-} > C_{2} H_{5} O^{-} > \left(OH\right)^{-}\)

4. \(C_{2} H_{5} O^{-} > \left(OH\right)^{-} > CH \equiv C^{-}\)

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, as per the expression:

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.
Henderson's equation is used to determine pH of buffer mixtures of different 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? 
[Given: K_a of HCN = \(1 \times 10^{-10}\)]

1. 0.009

2. 0.09

3. 0.9

4. Buffer solution cannot formed