When 10 ml of 0.1 M acetic acid (pK_a=5.0) is titrated against 10 ml of 0.1 M
ammonia solution (pK_b=5.0), the pH at equivalence point will be:

The salt solution that is basic in nature is: 

1. Ammonium chloride.

2. Ammonium sulphate.

3. Ammonium nitrate.

4. Sodium acetate.

The degree of hydrolysis of the salt is independent of the concentration of a solution: 
1. NH₄CI
2. NH₄CN
3. (NH₄)₂SO₄
4. All of the above

The salt that gives a neutral solution in water is: 

1. KBr

2. NH₄NO₃

3. NaCN

4. Rb₂(CO₃)

The ionization constant of chloroacetic acid is 1.35 × 10^–3. The pH of a 0.1 M acid solution will be:

The percentage degree of hydrolysis of a salt of weak acid (HA) and weak base (BOH) in its 0.1 M solution is found to be 10%. If the molarity of the solution is 0.05 M, the percentage hydrolysis of the salt should be -

1. 5 %

2. 10 %

3. 20 %

4. None of the above

If the pH of 0.1 M NaCN solution is 11, then the percentage hydrolysis will be:

1. 0.1%

2. 1%

3. 0.01%

4. 10%

The ionization constant of nitrous acid is 4.5 × 10^–4.  The pH of a 0.04 M sodium nitrite solution will be:

1. 2.56

2. 6.14

3. 7.96

4. 11.74

A 0.02 M solution of pyridinium hydrochloride has a pH of 3.44.
The ionization constant of pyridine will be:

1. $6.6$ $\times$ ${10}^{-6}$

2. $1.51$ $\hspace{0.17em}\times$ ${10}^{-9}$

^{3. $2.36$ $\times$ ${10}^{-7}$}

4. $3.43$ $\times \hspace{0.17em}{10}^{-9}$