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:

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.	None of the above

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

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

The pH of 10^-6 M CH₃COOH will be:
(Given: k_a of CH₃COOH = 1.8 \(\times\) 10^-5  & log 4.24 =0.63)

1. 5.37

2. 7.0

3. Slightly more than 6

4. 6.95

What can be the active mass of CaCO₃ if 10 grams of CaCO₃ are taken in a one-liter container?
[Given: Molecular weight of CaCO₃ = 100]

1. 0.1 

2. 1 

3. 0.01 

4. 10