Which composition will make the basic buffer?

An acidic buffer cannot be formed by which of the following combinations is:

1. ${\mathrm{HClO}}_4$ $\mathrm{and}$ ${\mathrm{NaClO}}_4$
2. ${\mathrm{CH}}_3\mathrm{COOH}$ $\mathrm{and}$ ${\mathrm{CH}}_3\mathrm{COONa}$
3. ${\mathrm{H}}_2{\mathrm{CO}}_3$ $\mathrm{and}$ ${\mathrm{Na}}_2{\mathrm{CO}}_3$
4. ${\mathrm{H}}_3{\mathrm{PO}}_4$ $\mathrm{and}$ ${\mathrm{Na}}_3{\mathrm{PO}}_4$

Buffer solutions have constant acidity and alkalinity because:
 

A buffer solution is prepared in which the concentration of NH₃ is 0.30 M and the concentration of  $N{H}_4^{+ }$is 0.20 M. If the equilibrium constant, K_b for NH₃ equals 1.8×10^-5, then what is the pH of this solution? 
(log 1.8 = 0.25; log 0.67 = -0.176)

1.  9.43
2.  11.72
3.  8.73
4.  9.08

What is [H⁺] in mol/L of a solution that is 0.20 M in CH₃COONa and 0.10 M in CH₃COOH ?(K_a  for CH₃COOH = 1.8 x 10^-5):
1. $3.5\times {10}^{-4}$
2. $1.1\times {10}^{-5}$
3. $1.8\times {10}^{-5}$
4. $9.0\times {10}^{-6}$

In a buffer solution containing an equal concentration of B^- and HB, the K_b for B^- is 10^-10. pH of the buffer solution is:

The following pair constitutes a buffer is:

1. $HN{O}_2$ $and$ $NaN{O}_2$

2. $NaOH$ $and$ $NaCl$

3. $HN{O}_3$ $and$ $N{H}_{4}N{O}_3$

4. $HCl$ $and$ $KCl$

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

4. All of the above.