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. $\log \frac{\left[{\mathrm{In}}^{-}\right]}{\left[\mathrm{HIn}\right]}=\; pH\; +\hspace{0.17em}{\mathrm{pK}}_{\mathrm{In<mspace\; width="1em"></mspace>}}$

4. None of the above

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? 
(k_a of HCN = \(1 \times 10^{-10}\))

1. 0.009

2. 0.09

3. 0.9

4. Buffer solution cannot formed